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GIFT  OF 
Harry  East  Miller 


A  LABORATORY  MANUAL 

OF 

General  Chemistry 

For  Use  in  Colleges 


BY 

WILLIAM  CMBRAY, 

Professor  of  Chemistry  in  the 
University  of  California 


AND 


WENDELL   M.  LATIMER 

Instructor  of  Chemistry  in  the 
University  of  California 


A  LABORATORY  MANUAL 

OF  | 

General  Chemistry 

For  Use  in  Colleges 


BY 
WILLIAM  C.  BRAY, 

Professor  of  Chemistry  in  the 
University  of  California 


AND 


WENDELL   M.   LATIMER 

Instructor  of  Chemistry  in  the 
University  of  California 


Copyright,  1921,  By  William  C.  Bray  and  Wendell  M.  Latimer 
Price,  Fifty  Cents 


LEDERER,  STREET  &  ZEUS  Co.,  Publishers 
BERKELEY,  CALIFORNIA 

1921 


PREFACE 


The  present  laboratory  manual  has  been  prepared  primarily  for  the  use  of 
students  in  general  inorganic  chemistry  in  the  University  of  California.  These 
students  have  usually  had  a  year  in  elementary  chemistry  in  high  school,  and 
many  of  them  will  take  no  further  work  in  chemistry.  No  distinction  is  made 
between  students  on  the  basis  of  the  various  curricula  which  they  are  following, 
as  we  believe  that  a  course  in  the  fundamentals  of  general  chemistry  is  equally 
suitable  for  all  students. 

The  laboratory  and  lecture  work  are  correlated  as  closely  as  possible.  In  the 
present  manual  page  references  are  given  to  Professor  Joel  H.  Hildebrand's 
"Principles  of  Chemistry/'  MacMillan,  1918,  the  reference  book  written  for  the 
course. 

The  laboratory  work  is  a  study  of  chemical  principles,  rather  than  a  presen- 
tation of  descriptive  material.  It  is  hoped  that  the  division  of  the  manual  into 
Sections,  and  the  statements  in  the  first  paragraphs  of  the  various  Assignments, 
will  materially  assist  the  student  in  recognizing  the  relation  between  the 
experimental  details  and  the  principles  involved. 

When  the  course  extends  over  two  terms,  as  at  the  University  of  California,  a 
satisfactory  division  is  to  take  Sections  I  to  III  in  the  first  term,  though  in  some 
cases  it  may  be  possible  also  to  begin  the  first  Assignment  on  Qualitative  Analysis. 
It  is  recommended  that  the  Assignments  in  the  last  two  Sections  be  taken  in  the 
order  noted  in  the  text. 

The  following  editions  of  the  manual  have  been  printed  :  Laboratory  Directions 
in  Chemistry  1A,  edited  by  William  C.  Bray,  1915;  21  Assignments.  A  Labora- 
tory Manual  of  General  Chemistry,  William  C.  Bray  and  Ludwig  Rosen  stein, 
1916;  26  Assignments.  The,  6  s^rne,  .  revised,  by.  William  C.  Bray,  1917;  31 
Assignments;  reprinted  1918,  \9i9,l92^J  frfiS  present  manual  contains  5  Sections 
with  a  total  of  30  Assignments,  and  is  an*  almost  complete  revision  of  the  1917 

manual.  A  :  .:::  :".i  :Y  '•:  :    :/•.'"": 


In  the  development  of  this  manual  from  1912  to  the  present  time  a  great  deal 
has  been  contributed  by  the  instructors  in  the  course.  We  wish  especially  to 
acknowledge  our  indebtedness  to  Professors  G.  N.  Lewis,  J.  H.  Hildebrand, 
Edward  Booth  and  E.  D.  Eastman  and  to  Doctors  Ludwig  Rosenstein  and 
W.  L.  Argo. 

WILLIAM  C.  BRAY, 
WENDELL  M.  LATIMER. 


June,  1921. 


* 

a 


TABLE  OF  CONTENTS 


Pago 
Note  to  Students 5 

Special  Laboratory  Directions 5 

List  of  Apparatus 6 

SECTION  I.    WEIGHT  RELATIONS  IN 
CHEMICAL  REACTIONS. 

Assignment  1.     A  Chemical  Reaction:     The  Synthesis  of  a  Sulfide 

of  Copper 7 

Notes  on  Glass   Manipulation 8 

Assignment  2.     The   Relation  between  the  Mass  and   Volume  of 

Gases:    The  Determination  of  the  Volume  of  a  Mol  of  Oxygen.. 9 

Assignment  3.     The  Reaction  between  Certain  Metals  and  Hydrochloric 

Acid 11 

Assignment  4.    The  Analysis  of  Copper  Oxide 13 

Assignment  5.    The  Reaction  between  an  Acid  and  a  Base  in  Solution. 

Concentration  in  Terms  of  Mols  per  Liter 15 

Assignment  6.     The  Titration  of  Solutions  of  Acids  and  Bases :  An 

Illustration    of   Volumetric    Analysis 17 

Assignment  7.    Volumetric  Analysis,  Continued :  The  Determination  of 

the  Equivalent  Weight  of  an  Unknown  Acid 19 

SECTION  II.     IONIC  THEORY. 

RAPID  REVERSIBLE  REACTIONS  AND  EQUILIBRIUM 
Assignment  21.     Solutions  of  Strong  Electrolytes 21 

Assignment  22.     Strong  and  Weak  Acids.     The  Use  of  Indicators 

to  Measure  Hydrogen  Ion  Concentration 23 

Assignment  23.     Strong  and  Weak  Bases.     The  Use  of  Indicators 

to  Measure  Hydroxide  Ion  Concentration 26 

Assignment  24.    Rapid  Reversible  Reactions  and  Equilibrium 27 

Assignment  25.    The  Reversibility  of  Neutralization  Reactions.    Hydrolysis....  30 

SECTION  III.    REACTIONS  OF  IONS 

Assignment  31.     The  Properties  of  Sodium,  Potassium  and  Ammonium 

Ions.     Tests  for  Chloride,  Sulfate  and  Nitrate  Ions 33 

Assignment  32.     Calcium  Ion 36 

Assignment  33.     Carbonate  Ion,  Bicarbonate  Ion  and  Carbonic  Acid 38 

Assignment  34.     Sulfates,  Chlorides  and  Nitrates  of  Copper,  Silver  and 

Zinc 42 

Assignment  35.     Hydroxides  of  Copper,  Silver  and  Zinc 44 

Assignment  36.     Complex  Ions  of  Copper,  Silver  and  Zinc  with  Ammonia 46 


Page 
Assignment  37.     Carbonates  and  Sulfides  of  Copper,  Silver  and  Zinc 48 

Assignment  38.     Review  of  the  Chemistry  of  Positive  Ions  Already 

Considered -   50 

SECTION  IV.     REACTIONS  OF  IONS,   CONTINUED 

Assignment  41.    Oxidation  and  Reduction.    Replacement  Reactions. 

Electrical  Cells 52 

Assignment  42.    Oxidation  of  Metals  to  their  Ions.    Table  of  Oxidizing 

and  Reducing  Agents 54 

Assignment  43.  Ferrous  and  Ferric  Ions 56 

Assignment  44.  Mercurous  and  Mercuric  Ions 58 

Assignment  45.  Lead  Ion,  Chromate  Ion 61 

Assignment  46.  Stannous  and  Stannic  Ions.  Amphoteric  Sulfides 61 

Assignment  47.  Ions  of  Aluminum  and  of  Chromium.     Peroxides 62 

SECTION  V.    QUALITATIVE  ANALYSIS 

Assignment  51.    The  Development  of  a  Scheme  of  Analysis  for  a 

Limited  Number  of  Positive  Ions 65 

(To  follow  Assignment  38) 

Assignment  52.    The  Standard  Scheme  of  Analysis.     Methods  of 

Dissolving  Difficulty   Soluble   Substances 67 

(To  follow  Assignment  44) 

Assignment  53.     Lead,  Tin,  Aluminum,  Chromium  and  Barium  in  the 

Scheme  of  Analysis 70 

(To  follow  Assignment  47) 


NOTE  TO  STUDENTS 

1.  Decide  what  is  the  real  purpose  of  each  Assignment.     Before  beginning 
the  experimental  work  and  preferably  before  coming  to  the  laboratory,  read  the 
first  paragraphs  of  the  Assignment,  study  carefully  the  References,  and  review 
earlier,  related  work.     In  general  look  for  the  connection  between  the  lectures 
and  laboratory  work  and  between  each  Assignment  and  the  preceding  ones. 

2.  Master   each   idea   before   proceeding   to    the   next   one.      Form   the   habit 
of  at  once  consulting  the  instructor  whenever  you  are  not  certain  of  the  cor- 
rectness of  your  answer  to  questions  and  of  your  conclusions  from  the  experi- 
ments.    The  frequent  short  written  examinations  will  be  of  great  assistance  to 
you  in  deciding  whether  or  not  you  have  really  understood  the  work. 

3.  An  average  student  who   has  understood   the    earlier  work   can   complete 
an  Assignment  in  the  regular  laboratory  time  allotted  to  it.    Students  who  can- 
not finish  in  the  stated  time,  announced  by  the  instructor,  must  consider  them- 
selves behind  the  class,  and  should  plan  immediately  to  do  extra  work  at  home 
and    in    the    laboratory.      For    their    convenience    the    building    is    open    from 
8:00  A.M.  to  4:30  P.M.   (Saturdays  8:00  to   12:00). 

4.  Your   success   depends    upon    your    own    efforts.      If   you    are    in    serious 
difficulty  then  there  is  something  wrong  with  your  methods.     The   instructor 
can  assist  you  in  finding  out  what  is  wrong,  but  he  cannot  do  your  work  for 
you.     No  effort  on  his  part  can  make  up  for  lack  of  initiative  on  your  part, 
failure   to   assume   the    responsibility   of   mastering   each    idea,   or   inability   to 
improve  your  method  of  doing  the  work. 

SPECIAL  LABORATORY  DIRECTIONS 

5.  A  laboratory  note  book  about  6^   inches  wide,   opening  at  the  side,  and 
not  loose-leaved,  is  recommended.   This  book  and  the  laboratory  manual  will  be 
needed  at  each  meeting  of  the  laboratory  section,  including  the  first  one. 

6.  The    recording   of   experiments,    observations   and   conclusions    at   once   in 
the  note  book  is  an  essential  part  of  the  laboratory  work.     Entries  made  from 
memory  or  from  memoranda  on  scraps  of  paper  are  not  records  of  the  experi- 
mental work.     The  value  of  the  original  record  is  improved:  by  dating  each 
day's  notes;  by  numbering  the  pages  of  the  note  book;  by  never  erasing  an 
entry  or  tearing  out  a  page;  by  leaving  space  for  additions  and  corrections; 
by  adopting  a  plan   of  distinguishing  between  the  record   of  the  experiments 
actually  performed,  and  the  other  entries  such  as  answers  to  questions,  solu- 
tions of  problems,  etc. ;  by  making  all  calculations  neatly  at  the  bottom  or  side 
of  the  page ;  and  by  writing  entries  in  such  a  way  that  they  will  be  easily  under- 
stood when  the  work  is  reviewed.     The  descriptions  of  the  experiments  per- 
formed should  be  very  brief  when  detailed  directions  are  given  in  the  manual, 
but  must  be  complete  when,  as  in  the  later  work,  experiments  are  devised  by 
the  student.     A  passing  grade  in  the  laboratory  work  will  not  be  given  unless 
the  experiments  have  been  completed  and  the  results  properly  recorded  in  the 
note   book. 

7.  The  laboratory  desk  must  be  kept  neat  and  dry.     An  old  towel  should  be 
used  for  cleaning  the  desk  top  and  another  towel  should  be  kept  clean  for  use 
on  apparatus.     When  cleaning  apparatus  use  tap  water  and  a  brush  to  remove 
all  visible  dirt  and  rinse  finally  with  a  little  distilled  water.     Before  leaving  the 
laboratory  the  apparatus  should  be  locked  up  in  the  desk. 

8.  The  wash-bottle  should  only  be  used  to  hold  distilled  water.     Before  using 
sterilize   the    mouth-piece   by    boiling    in    water,    and   never    lend   or   borrow    a 
wash-bottle. 


9.  The  contamination  of  laboratory  reagents  can  be  avoided  by  keeping  each 
stopper  clean  and  replacing  it  at  once  in  the  proper  bottle,  and  by  never  pouring 
anything  back  into  a  reagent  bottle. 

10.  Experiments   which   give   rise   to   disagreeable   or   dangerous   fumes   must 
always  be  performed  out  of  doors  or  in  a  fume-closet. 

11.  At  the  first  meeting   of  each  laboratory   section   the  instructor  will  dis- 
tribute the  desk  keys  together  with  lists  of  apparatus  similar  to  the  one  given 
below.     Check  the  apparatus  in  the  locker,  exchange  damaged  articles  at  the 
office,  sign  the  list  of  apparatus  (surname  first)  and  return  it  to  the  instructor. 
Begin  work  on  Assignment  I. 

LIST  OF  APPARATUS 

1.  Regular  equipment  of  each  locker.  Additional  articles  may  be  obtained 
at  the  office  by  filling  out  an  "order  slip"  and  signing  your  name  and  locker 
number.  Whenever  any  article  is  returned  to  the  office  sign  a  "return  slip."  At 
the  end  of  the  term  the  locker  must  contain  the  same  amount  of  apparatus,  no 
more  and  no  less;  the  locker  must  be  clean;  the  apparatus  must  be  clean  and 
dry,  and  in  good  condition;  glass  stoppers  must  fit,  and  be  protected  by  the 
insertion  of  a  piece  of  paper. 


1  Key. 

5  Beakers,  100  cc.,  150  cc.,  250  cc., 

400  cc.,  600  cc. 
5  Reagent  Bottles. 

2  Sample  Bottles,  50  cc. 

1   Graduated  Cylinder,  50  cc. 
4  Flasks,  500  cc.,  250  cc.,  and  two 
125  cc. 

1  Wash-bottle,   equipped   with   glass 

tubing  and  rubber  stopper. 

2  Funnels. 

2  Blue  Glasses. 

2  Glass  Rods,  12  cm.  and  18  cm. 
30  cm.  Glass  tubing.* 
12  Test-tubes. 

1  Watch  Glass. 


1  Casserole. 

1  Crucible,  with  cover. 

2  Evaporating  Dishes. 
1  Crucible  Tongs. 

1   Bunsen  Burner,  with  rubber  tubing. 

1  Iron  Wire.* 

1  Wire  Gauze.* 

1  Triangle. 

1  Test-tube  Brush.* 

1  Test-tube  Holder. 

1  Test-tube  Rack. 

1   Package  Filter  Paper.* 

1  Rule. 

2  Towels. 

Litmus  Paper*  in  a  bottle. 


2.  The  following  additional  articles  may  be  obtained  at  the  office : 

(a)  By  signing  the   regular  order  slips.      Small   short-stemmed    funnels ;   glass 
flasks,  50  cc. ;  matches  * ;  corks ;  rubber  stoppers. 

(b)  By  signing  "temporary  order  ships."     Special  apparatus   for  Assignments 
2  and  4;  burettes,  with  clamps  and  pinch-cocks;  graduated  cylinders,  10  cc. 
and  250  cc. ;   thermometers ;  paraffin.     These  articles   should   be   returned 
when  possible  during  the  same  laboratory  period. 


*Not  returnable.     At  the  end  of  the  first  term  students  should  retain  these  articles  for 
use  in  the  second  term. 


SECTION  i  v /.\ \fj \  \-  v  K  •-  / 

WEIGHT  RELATIONS  IN  CHEMICAL  REACTIONS 


ASSIGNMENT  1 

A  CHEMICAL  REACTION  :    THE  SYNTHESIS  OF  A 
SULPHIDE  OF  COPPER 


References.     Hildebrand,  Principles  of  Chemistry,  Chapter  I,  and  pages  40-43. 

1.  In  Assignment  1  we  shall  study  quantitatively  a  chemical  reaction  in  which 
two  elements,  *  a  metal  and  a  non-metal  unite  to  form  a  pure  compound.     The 
experiment  consists  in  determining  the  weight  of  the  compound  that  is  formed 
from  a  weighed  amount  of  copper  when  heated  with  excess  of  sulfur.     From 
these  experimental  data,  and  the  atomic  weights  of  the  two  elements,  the  relative 
number   of   atoms  of   copper  and   sulfur  in  the   compound   will   be   calculated. 
Questions.    If  3.04  grams  of  a  certain  metal,  when  burned  in  oxygen,  yield  5.04  g. 
of  a  pure  compound  of  the  metal  and  oxygen,  what  weight  of  oxygen  will  combine 
with  1.00  g.  of  this  metal?     What  additional  information  is  necessary  before  the 
relative  number  of  atoms  of  the  two  elements  in  the  compound  can  be  calculated  ? 

2.  Experiment.    Support  a  clean  porcelain  crucible,  with  a  cover,  on  a  triangle 
and  heat  with  the  colorless  flame  of  a  bunsen  burner  to  lowr  redness.     Let  the 
crucible  cool  about  10  minutes,  and  weigh  it,  with  the  cover,  to  10  milligrams. 
Note.     Do  not  make  any  weighings  until  instructions  in  the  use  of  the  balance 
have  been  given. 

3.  While  the  crucible  is  cooling  obtain  from  the  shelf  a  clean  piece  of  copper 
wire,  weighing  about  1  gram,  and  weigh  it  to  10  mg. 

4.  Place  the  copper  in  the  weighed  crucible  and  add  enough  powdered  sulfur 
to  cover  the  copper.     Place  the  cover  on  the  crucible  and  heat  gently   (with  a 
small  flame)  until  the  sulfur  ceases  to  burn  at  the  edges  of  the  cover,  but  do  not 
remove  the  cover  while  the  crucible  is  hot.     Then  heat  more  strongly  until  the 
bottom  of  the   crucible  just  becomes  dull  red.     Again  allow  to  cool  about  10 
minutes  and  weigh. 

5.  Carefully  remove  the  cover  and  note  the  appearance  of  the  contents  of  the 
crucible,  but  do  not  touch  the  substance.     If  there  is  any  free  sulfur  on  the  cover 
or  the  wall  of  the  crucible,  replace  the  cover,  heat  the  crucible  and  cover,  and 
weigh  again.    Check  the  accuracy  of  the  final  weight  by  adding  a  small  quantity 
of  sulfur  and  repeating  the  experiment ;    continue  until  two  consecutive  results 
agree  within  10  mg.    At  the  end  of  the  experiment  remove  the  substance  formed, 
break  it  and  describe  its  properties.     Clean  the  crucible  with  hot  nitric  acid  in 
a   porcelain   dish,   wash   with   distilled   water,    dry   by   heating,   and   check   the 
original  weight. 

6.  Questions.     What  conclusions  can  you  draw  from  each  of  the  following 
observations:     (a)    the  properties   of  the  product  are  different   from   those  of 
either  copper  or  sulfur;    (b)   the  product  appears  to  be  homogeneous  and  its 
weight  is  greater  than  that  of  the  copper  used?     What  additional  evidence  is 
necessary  to  prove  that  the  product  is  a  pure  substance  and  not  a  solid  solution? 

7.  Calculations.     Summarize  your  experimental  results  and  make  the  calcula- 
tions necessary  to  complete  a  table  similar  to  the  following: 

*  It  is   suggested  that  the   student   form  the  habit  of  writing  out  the  meaning  of  each 
italicized  word  in  the  text  and  of  giving  an  example  whenever  possible. 

[7] 


(a)  Weight  of  crucible 

(b)  Weight  of  copper 

(c)  Weight  of  crucible  and  product 

(d)  Weight  of  product 

(e)  Difference  between  (d)  and  (b) 


First          Second          Value 
Weighing    Weighing       Chosen 


Calculate  what  the  increase  in  weight  (e)  would  have  been  if  one  gram-atom 
of  copper  had  been  used  in  the  experiment,  and  enter  in  the  table  as  line  (/)  ; 
show  this  result  to  your  instructor  at  once.  How  does  this  number  compare  with 
the  atomic  weight  of  sulphur?  How  many  gram  atoms  of  sulfur  have  combined 
with  one  gram  atom  of  copper?  What,  then,  is  the  simplest  formula  of  the 
substance  formed?  What  is  the  corresponding  molecular  weight?  Write  the 
equation  for  the  reaction,  and  write  out  in  words  what  this  equation  means,  in 
terms  of  (a)  atoms  and  molecules,  (b)  gram  atoms  and  mols,  (c)  grams,  and 
(d)  pounds. 

8.  Problems.  (/)  In  order  to  determine  the  effect  of  a  small  error  in 
weighing  the  copper  repeat  your  calculations,  Paragraph  7,  using  for  the  weight 
of  copper  a  value  10  mg.  greater  than  your  experimental  value.  What  per  cent 
of  the  weight  of  copper  is  10  mg.  ?  This  would  be  the  percentage  error  in  the 
weight  of  copper  if  a  10  mg.  error  in  weighing  had  been  made.  What  is  the 
corresponding  percentage  error  in  your  value  for  the  weight  of  sulfur  that 
would  combine  with  one  gram  atom  of  copper? 

(2)  Calculate  the  percentage  composition  of  the  copper  sulfide  formed  : 
(a)  from  your  experimental  data,  and  (b)  from  the  formula  and  the  atomic 
weights  of  copper  and  sulfur.  Compare  the  results. 

(j)  A  sulfide  of  iron  contains  53.8%  iron.  What  is  the  formula?  (In  solving 
this  problem  consider  one  gram  atom  of  iron  and  one  gram  atom  of  sulfur  as 
the  fundamental  units  for  iron  and  sulfur.  Calculate  first  the  weight  of  sulfur 
and  then  the  number  of  gram  atoms  of  sulfur  combined  with  one  or  more 
gram  atoms  of  iron.) 

(4)  The  formulas  of  hydrogen  sulfide  and  of  ferrous  sulfide  are  HL.S  and 
FeS,  respectively.  What  are  their  molecular  weights.  What  weight  of  sulfur  is 
contained  in  one  mol  of  hydrogen  sulfide?  In  one  mol  of  ferrous  sulfide?  What 
weight  of  hydrogen  sulfide  could  be  made  from  one  mol  of  ferrous  sulfide? 


NOTES  ON  GLASS  MANIPULATION 

To  bend  a  piece  of  ordinary  glass  tubing,  hold  it  with  both  hands  in  a 
fan-shaped  gas  flame  and  rotate  it  slowly  between  the  thumb  and  fingers  until 
a  2]/2.  to  3  inch  portion  is  uniformly  heated  and  is  soft  enough  to  be  bent  to  the 
proper  angle.  Set  it  aside  to  cool ;  glass  will  remain  hot  enough  to  burn  the 
hand  for  some  time  after  it  no  longer  appears  to  be  hot. 

To  cut  glass  tubing,  scratch  it  with  a  file  at  the  proper  place,  grasp  it  firmly 
on  each  side  of  this  mark  (protecting  the  hands  with  a  cloth),  and  bend  the 
tube  away  from  the  mark. 

Always  remove  the  sharp  edges  of  freshly  cut  glass  at  once  with  a  file,  or  by 
heating  in  a  gas  flame. 

To  draw  down  a  piece  of  tubing  to  a  capillary,  heat  a  portion  about  1  inch 
long  in  an  ordinary  gas  flame  to  a  higher  temperature  than  was  necessary  in 
bending  the  tubing.  Hold  the  tube  with  both  hands  and  rotate  it  to  ensure 
uniform  heating  and  prevent  the  hot  portion  from  sagging.  Withdraw  from  the 
flame  and  draw  apart  slowly  to  obtain  a  thick-walled  capillary. 

[81 


ASSIGNMENT  2 

THE  RELATION  BETWEEN  THE  MASS  AND  VOLUME  OF  GASES:    THE 
DETERMINATION  OF  THE  VOLUME  OF  A  MOL  OF  OXYGEN 

References.    Hildebrand,  Chapter  II,  and  pages  52  and  57 

1.  It  is  often  necessary  to  know  the  volume  of  a  given  mass  of  a  substance,  or 
conversely  the  mass  of  a  given  volume.     While  in  the  case  of  a  solid  or  liquid 
the    relation    between   the   mass   and   the  volume  must   be   determined   for   the 
particular  substance,  the  problem  is  simplified  when  we  are  dealing  with  a  gas, 
since  a  mol  of  every  gas  occupies  nearly  the  same  volume  under  similar  condi- 
tions.     In   this   Assignment   we   shall   determine,    under   definite   conditions   of 
temperature  and  pressure,  the  volume  of  a  known  weight  of  oxygen  and  calculate 
the  volume  of  one  mol  at  standard  conditions.     Questions.     What  information 
is  needed  before  you  can  calculate  the  weight  of  10  cc.  of  mercury?     A  given 
solution  of  sodium  chloride  in  water  contains  25.0  percent  sodium  chloride  and  the 
density  of  the  solution  is  1.19  g.  per  cc. ;   what  volume  of  solution  in  cc.,  and  in 
liters,  contains  100  grams  of  sodium  chloride? 

2.  When  solid  potassium  chlorate  is  strongly  heated  it  decomposes  with  the 
evolution  of  oxygen,  and  the  loss  in  weight  gives  the  weight  of  oxygen  evolved. 
The  volume  of  the  oxygen  is  determined  by  measuring  the  volume  of  water 
displaced  by  an  equal  volume  of  oxygen.    When  pure  potassium  chlorate  is  used 
it  must  be  heated  in  a  hard  glass  (difficulty  fusible)  test-tube.     The  potassium 
chlorate   decomposes   more   readily  and  at  a  lower  temperature  when  a  small 
quantity  of  manganese  dioxide  is  present.     The  hard-glass  test-tube  may  then 
be  replaced  by  a  heavy-walled  test-tube  of  ordinary  easily-fusible  glass ;   but  care 
must  be  taken  not  to  heat  the  latter  tube  to  a  higher  temperature  than  is  necessary 
for  the  reaction.     The  manganese  dioxide  is  a  catalyst  in  this  reaction,  and  all 
of  it  may  be  recovered  after  the  potassium  chlorate  has  been  decomposed  into 
potassium  chloride  and  oxygen. 

3.  Two  students  working  together  obtain  from  the  office  a  heavy-walled  glass 
tube,  a  rubber  stopper,  rubber  tube,  pinch-cock,  clamp  and  small  tube  containing 
about  5  grams  potassium  chlorate.    The  yellow  order  slip  for  "special"  apparatus 
for  Assignment  2  should  be  signed  by  both  students.     The  apparatus  should  be 
returned  as  soon  as  the  experiment  is  finished.    Each  student  must  keep  a  complete 
record  of  the  experiment  in  his  notebook. 

4.  Experiment.     Set  up  the  apparatus  according  to  the  accompanying  diagram. 
Since  the  apparatus  must  be  gas  tight,  glass  tubing  and  stoppers  must  be  care- 
fully fitted.     If  your  tubes  and  stoppers  do  not  fit,  exchange  them  at  the  office. 
Do  not  use  the  glass  tubing  of  your  wash  bottle.     Directions  for  bending  glass 
tubing  are  given  on  page  8. 

5.  Place  in  the  heavy-walled  test-tube  about  5  grams  potassium  chlorate.    Add 
about  50  mg.  manganese  dioxide,  estimating  the  amount  by  comparison  with  the 

sample  in  the  laboratory;  mix  it  with  the 
potassium  chlorate,  by  jarring  the  tube;  and 
wipe  off  any  powder  that  is  on  the  outside  of 
the  tube  or  on  the  inside  near  its  mouth. 
Assemble  the  apparatus  as  before. 

6.  Heat  the  tube  gently  with  a  moving  gas 
flame,  leaving  the  pinch-cock  open  on  the 
rubber  tube  outlet.  Moisture  will  appear  on 
the  walls  of  the  test-tube,  which  shows  that 
the  potassium  chlorate  and  manganese  dioxide 
were  not  perfectly  dry.  Gradually  warm  the 
tube  to  within  about  one  inch  of  the 


stopper  and  at  the  same  time  heat  the  potassium  chlorate  until  gas 
evolution  begins  and  some  water  passes  over  into  the  beaker.  Drive  out 
the  moisture  with  the  oxygen  by  carefully  heating  the  walls  of  the  tube, 
but  do  not  scorch  the  rubber  stopper.  When  50  to  100  cc.  water  have  passed  into 
the  beaker,  allow  the  apparatus  to  cool;  close  the  pinch-cock  near  the  outer  end 
of  the  rubber  tube,  which  should  now  be  filled  with  water.  Disconnect  the 
chlorate  tube,  and  weigh  it  with  the  dry  material  inside  to  10  mg. 

7.  Immediately  replace  the  tube  in  its  proper  position,  open  the  pinch-cock 
while  the  end  of  the  delivery  tube  is  under  the  water,  raise  the  beaker  until  the 
surfaces  of  the  water  inside  and  outside  the  flask  are  at  the  same  level,  close  the 
pinch-cock  again,  place  a  dry  beaker  under  the  delivery  tube,  and  open  the  pinch- 
cock.     If  the  water  continues  to  siphon  into  the  beaker  your  apparatus  is  not  gas 
tight  and  must  be  rebuilt.     Again  heat  the  tube  gradually  until  gas  evolution 
begins,    and    continue    to   heat   the   potassium    chlorate    carefully    and    not   too 
strongly  until  from  250  to  300  cc.  water  have  been  forced  over  into  the  beaker. 
Allow  the  tube  to  cool,  equalize  the  level  of  the  water  in  the  beaker  and  the  flask, 
and  then  close  the  pinch-cock  on  the  siphon  tube.    By  means  of  a  250  cc.  graduated 
cylinder  measure  the  amount  of  water  which  the  oxygen  has  forced  out  of  the 
flask.      Finally    weigh    carefully    the   hard    glass    tube    containing   the    partially 
decomposed  chlorate.     Question.    Why  is  it  necessary  to  cool  the  test-tube  and  to 
have  the  water  in  the  beaker  and   flask  at  the   same  level   before   closing  the 
pinch-cock? 

8.  Repeat  the  experiment,  Paragraph  7.    The  three  weighings  and  two  volume 
measurements  give   two   independent   sets   of  experimental  data.     Compare  the 
results  of  your  two  experiments  by  preparing  a  table  which  will  show  for  each 
experiment :    the  weight  of  oxygen,  the  corresponding  volume  of  water  displaced, 
the  data  referred  to  in  10  to  11  below,  and  the  results  of  the  calculations  (12). 

9.  Clean  the  test-tube  by  placing  water  in  it  and  shaking.     The  manganese 
dioxide  is  difficultly  soluble  in  water,  while  both  potassium  chlorate  and  potas- 
sium chloride  dissolve  readily.     Suggest  a  method  of  recovering  the  manganese 
dioxide    and    obtaining    a    mixture    of    dry    potassium    chloride    and    potassium 
chlorate  practically  free  from  manganese  dioxide. 

10.  To   make   the   calculations  it   will   be  necessary  to  know  the  barometric 
pressure  at  the  time  you  perform  each  experiment,  and  the  temperature  of  the 
water  in  the  flask.    The  temperature  of  the  water  may  be  assumed  to  be  that  of 
the  room,  and  the  barometric  pressure  will  be  posted  on  the  blackboard.     Enter 
these  data  in  your  notebook  before  leaving  the  laboratory.     Below  is  given  a  table 
of  the  vapor-pressure  of  water  at  different  temperatures : 

Vapor  Pressure  of  Water. 

Temp.  °C.  Vapor  Pressure  Temp.  °C.  Vapor  Pressure 

14  1.2  cm.  mercury  24  2.2  cm.  mercury 

16  1.3     "  '       "  26  2.5     " 

18  1.5     "          "  28  2.8     " 

20  1.7     "         "  30  3.2     " 

22  2.0    "  32  3.5     " 

11.  Questions.    Assuming  that  levels  of  the  water  in  the  beaker  and  the  flask 
were  the  same  when  the  pinch-cock  was  closed,  what  was  the  total  pressure  of 
the  gas  in  the  flask?     What  was  the  partial  pressure  of  the  water- vapor?     Of 
the  oxygen? 

12.  Calculations.    From  each  of  your  two  sets  of  experimental  data,  by  means 
of  the  Gas  Laws,  calculate  the  volume  of  1  mol  (32  grams)  of  pure  oxygen  at 
1  atmosphere  pressure  and  O°  C.     Compare  your  results  with  the  value  given  in 
your  text  book.     Your  results  should  not  differ  from  this  value  by  more  than 

F  101 


5    percent    (check    your    calculations).     Show    your    tabulated    results    to    your 
instructor,  and  repeat  the  experiment  if  necessary. 

13.  Questions.     What  is  the  formula  of  the  oxygen  molecule?    What  value  is 
usually  accepted  as  a  close  approximation  for  the  volume  of  1  mol  of  gas  under 
standard  conditions?     Use  this  value  to  calculate  (a)  the  weight  of  a  liter  of 
hydrogen    chloride   gas,    HC1,    under    standard    conditions;    (b)    the   molecular 
weight  of  a  gas  whose  density  at  standard  conditions  is  known  to  be  0.001977 
g.  per  cc. 

14.  Write  an  equation  to  represent  the  decomposition  of  potassium  chlorate, 
KC1O3,   into  potassium   chloride,   KC1,   and   oxygen;    and  write  out  what  this 
equation  means  in  terms  of   (a)   molecules;    (b)   mols,  and   (c)   grams  of  the 
substance  involved. 

15.  Problems,     (i)  One  gram  of  potassium  chlorate  is  completely  decomposed 
into  potassium  chloride  and  oxygen.     Calculate  (a)   the  weight  of  oxygen  that 
could  be  obtained ;    (b)  the  volume  of  the  oxygen  (in  liters  and  in  cc.)  at  standard 
conditions,  and  (c)  the  volume  of  the  oxygen  at  27°  C.  and  750  mm.  mercury 
pressure. 

(2)  It  is  an  experimental  fact  that  2  volumes  of  carbon  monoxide  gas  react 
with  1  volume  of  oxygen  to  form  2  volumes  of  carbon  dioxide  gas.  Give  the 
reasoning  by  which,  from  this  result  you  can  conclude  that  the  molecule  of 
oxygen  contains  an  even  number  of  atoms. 

(5)  The  formulas  of  carbon  monoxide  and  carbon  dioxide  are  CO  and  CO2, 
respectively.  Write  the  equation  for  the  reaction  considered  in  the  preceding 
question,  and  interpret  it  in  terms  of  (a)  mols,  (b)  liters,  and  (c)  grams. 


ASSIGNMENT  3 
THE  REACTION  BETWEEN  CERTAIN  METALS  AND  HYDROCHLORIC  ACID 

References.     Hildebrand,  pages  84-86,  and  47-50. 

1.  Certain  metals,  aluminum,  zinc,  magnesium,  etc.,  react  with  a  solution  of 
an  acid,  with  evolution  of  hydrogen  gas  and  formation  of  a  salt  in  solution.     In 
this  assignment  we  shall  dissolve  a  definite  weight  of  a  metal  in  excess  of  hydro- 
chloric acid,  measure  under  definite  conditions  of  temperature  and  pressure  the 
volume  of  the  hydrogen  liberated,  and  calculate  the  weight  of  the  metal  that 
would   form   one  gram-atom   of  hydrogen.     From   this   result   and  the  atomic 
weight  of  the  metal  we  can  then  determine :    the  number  of  atoms  of  hydrogen 
formed  when  one  atom  of  the  metal  reacts  with  the  acid,  the  number  of  molecules 
of  acid   (HC1)   which  react  with  one  atom  of  metal,  and  the  formula  of  the 
chloride    of    the    metal    formed.      Questions.      What    experimental    facts    and 
reasoning  have  led  to  the  conclusion  that  the  formula  of  the  hydrogen  molecule 
is  H0?  If  the  valence  of  hydrogen  in  HC1  is  +  1  what  is  the  valence  of  the 
chlorine  in  this  compound? 

2.  The   instructor   will   supply   to   each   student   a   sample   of  a  metal   as  an 
"unknown."    Experiment.    Take  a  portion  of  the  metal  weighing  between  0.4  and 
0.5  g.    Clean  it,  if  necessary.    Weigh  to  5  mg. 

3.  Obtain  a  small  short-stemmed  funnel  at  the  office,  and  select  a  beaker  of 
such  size  that  the  funnel  when  placed  in  it  can  be  completely  covered  with  water. 
Place  the  weighed  metal  in  the  beaker,  place  the  inverted  funnel  over  it,  and 
pour  freshly  distilled  water  into  the  beaker  until  the  funnel  is  completely  covered. 
Note.    Tap  water  contains  a  relatively  large  amount  of  dissolved  air,  and  should 

[in 


not  be  used  in  this  experiment  unless  it  has  been  heated  to  boiling  to  expell  the 
greater  part  of  dissolved  air. 

4.  Pour  distilled  water  into  a  half  liter  flask  until  the  water  completely  fills  the 
flask.  Moisten  a  piece  of  filter  paper  slightly  larger  than  the  mouth  of  the  flask,  cover 
the  mouth  of  the  flask  with  paper,  taking  care  that  no  bubble  of  air  remains  below 
the  paper.     Invert  the  flask  (over  an  empty  vessel)  and  lower  it  into  the  beaker 
in  such  a  manner  that  the  stem  of  the  funnel  enters  the  neck  of  the  flask.     If 
a   bubble   of   air   enters   the   flask   repeat   this   operation.     The   apparatus   now 
consists  of  a  beaker  containing  a  funnel  inverted  over  the  metal,  and  a  flask 
filled  with  water  and  inverted  over  the  funnel.     Place  this  apparatus  in  a  large 
beaker  or  other  vessel,  to  prevent  the  water  from  overflowing  on  the  desk  during 
the  remainder  of  the  experiment. 

5.  Insert  a  thistle  tube  or  long-stemmed  funnel  into  the  water  so  that  the  lower 
end  touches  the  bottom  of  the  beaker  at  the  rim  of  the  inverted  funnel,  and 
through  it  pour  25  cc.  concentrated  hydrochloric  acid.     If  the  liquid  is  not  stirred 
the  concentrated  acid,  which  is  1.18  times  as  dense  as  water,  will  remain  for  some 
time  as  a  layer  at  the  bottom  of  the  beaker,  and  the  metal  will  be  dissolved 
rapidly.     If  all  the  water  in  the  inverted  flask  is  displaced  by  the  hydrogen  you 
have  used  too  much  metal  or  too  small  a  flask  and  must  begin  the  experiment 
over  again. 

6.  When   the  metal  has   all   dissolved    (except   a   few   dark-colored   flakes   of 
impurities   of  negligible  weight),   place   the   apparatus   in  a  large   basin  of  tap 
water  and  carefully  remove  the  beaker  and  funnel  without  allowing  any  air  to 
enter  the  inverted  flask.    Keep  the  flask  in  the  water  for  several  minutes  in  order 
that  it  may  be  at  the  same  temperature  as  the  water.     Then  raise  or  lower  the 
flask  until  the  level  inside  and  outside  the  flask  is  the  same.     (What  is  now  the 
pressure  of  the  gas  inside  the  flask?)    While  the  flask  is  in  this  position,  cover 
the  mouth  of  the  flask  with  the  palm  of  the  hand,  remove  the  flask  from  the  water 
and  invert  it.    While  the  gas  is  escaping,  test  to  prove  that  it  is  hydrogen. 

7.  Measure  the  volume  of  the  gas  which  was  contained  in  the  flask  by  filling 
the  flask  completely  with  water  and  observing  the  volume  needed.     Record  in 
your  notebook   the   barometric  pressure    (written  on  the  blackboard)    and   the 
temperature  of  the  water  in  which  the  flask  was  immersed. 

8.  You  now  have  the  weight  of  metal  taken,  and  the  volume,  at  a  definite 
temperature   and   pressure,   of   a   corresponding   amount   of   hydrogen   saturated 
with   water   vapor.     What   is   the  partial   pressure   of   the   water  vapor   at   the 
temperature  of  the  experiment?    What  was  the  partial  pressure  of  the  hydrogen 
in  the  flask? 

9.  Calculate  from  these  data : 

The  volume  at  standard  conditions  that  the  hydrogen  would  occupy  if  it 
were  dry. 

The  weight  of  the  hydrogen.  (Use  the  molecular  weight  2.016  and  the  volume 
of  1  mol  of  gas,  Assignment  2.) 

The  weight  of  metal  that  would  have  liberated  1  gram-atom  of  hydrogen. 
Report  this  value  to  your  instructor,  who  will  tell  you  the  name  of  the  metal  if 
your  result  is  correct  to  within  about  5%. 

By  means  of  the  atomic  weight  of  the  metal  calculate  the  number  of  atoms  of 
hydrogen  formed  when  1  atom  of  the  metal  dissolves  in  acid. 

10.  Questions.     How  many  molecules  of  HC1  react  with  1  atom  of  the  metal  ? 
Assuming  that  the  hydrogen  of  the  acid  is  replaced  by  the  metal,  what  is  the 
formula  of   the   chloride   formed?     What  is   the  valence  of  the   metal  in   this 
compound?     Write  the  equation  for  the  reaction. 

11.  Problems.      (i)    From   the   density   of   hydrogen   at   standard   conditions, 

[12.] 


0.00008987  g.  per  cc.,  calculate  the  actual  volume  of  1  mol  of  hydrogen.  What 
percentage  error  did  you  make  in  the  calculations  in  Paragraph  9  by  assuming 
the  value  22.40  liters. 

(2)  If  sulfuric  acid,  H2SO4,  had  been  used  in  the  above  experiment  instead  of 
hydrochloric  acid,  the  same  result  would  have  been  obtained  and  the  final  solution 
would  have  contained  a  sulfate  of  the  metal.  Write  the  equation  for  the  reaction. 

(j)  The  student  will  have  noted  that  in  the  first  three  Assignments  we  have 
assumed  a  knowledge  of  atomic  weights.  The  arbitrary  choice  of  the  unit 
0=16.00  should  present  no  difficulty,  but  it  is  often  not  clear  why  a  particular 
value  is  chosen  for  the  atomic  weight  of  an  element  rather  than  some  fraction  or 
multiple  of  this  value,  e.  g.,  in  the  case  of  chlorine  why  35.46  is  chosen  instead 
of  say  17.73  or  70.92.  To  illustrate  how  this  choice  is  made  on  the  basis  of 
the  experimentally  obtainable  quantities,  molecular  weight  and  percentage 
composition,  the  following  data  may  be  used.  (The  molecular  weights  given  in 
the  second  column  of  the  table  are  the  accurate  values,  and  not  approximate 
values  such  as  would  be  obtained  directly  from  the  weight  of  22.40  liters  of  gas 
reduced  to  standard  conditions.) 

No.  Grams  Chlo- 
Substance  Molecular  Weight     %  Chlorine       rine  in  i  Mol 

Chlorine 70.92  100 

Hydrogen  chloride 36.47  97.2 

Chlorine  oxide  (/) 86.92  81.6 

Chlorine  oxide  (?) 67.46  52.6 

Phosphorus  chloride  137.42  77.4 

Carbon  chloride  153.84  92.3 

Calculate  the  values  required  for  the  fourth  column  of  the  table.  What  value 
would  you  choose  for  the  atomic  weight  of  chlorine?  No  compound  of  chlorine 
has  even  been  made  which  contains  in  1  mol  less  than  35.46  grams  of  chlorine. 
How  many  atoms  of  chlorine  are  contained  in  a  molecule  of  each  of  the  six 
substances  listed  in  the  table? 


ASSIGNMENT  4 
THE  ANALYSIS  OF  COPPER  OXIDE 


Reference.    Hildebrand,  Chapter  III. 


1.  In  this  Assignment,  as  an  example  of  chemical  analysis,  we  shall  determine 
the  composition  of  an  oxide  of  copper.     The  analysis  will  be  made  by  heating  a 
weighed  portion  of  the  oxide  in  a  current  of  hydrogen  and  weighing  the  metallic 
copper  which  remains.     The  oxygen  of  the  oxide  unites  with  the  hydrogen  to 
form    steam.      As    in    Assignment    1,    we    shall    determine   the    formula   of    the 
compound  by  assuming  the  atomic  weights  of  copper  and  oxygen.     It  is  to  be 
noted,  however,  that  the  results  could  be  used  to  determine  the  relative  atomic 
weights  of   copper  and  oxygen  if  the   formula  of  the  compound  were  known. 
Our  experimental  data,  of  course,  will  not  be  sufficiently  accurate  to  make  worth 
while  the  calculation  of  the  atomic  weight  of  copper. 

2.  Experiment.    Two  students  may  work  together ;   both  should  sign  the  order 
slip  for  "special  apparatus  for  Assignment  4,"  which  consists  of  a  thick-walled 
hard  glass  test  tube  with  a  rubber  stopper  and  glass  tubes,  a  thistle  tube,  and 
two-holed  rubber  stopper,  a  clamp,  and  a  calcium  chloride  tube   (with  2  rubber 
stoppers,  2  glass  tubes,  2  rubber  tubes  and  2  short  glass  rods).     The  apparatus 
should  be  returned  as  soon  as  the  experiment  is  finished. 

3.  Set  up  a  ''hydrogen  generator"  by  fitting  your  half  liter  flask  with  a  thistle 
tube  extending  through  the   rubber  stopper  nearly  to  the  bottom  of  the  flask, 

[13] 


and  an  outlet  tube  bent  at  right  angles  (see  note  on  glass  manipulation).  Place 
in  the  flask  about  10  grams  of  zinc  and  cover  it  with  about  100  cc.  water.  To  the 
outlet  tube  attach  a  "drying  tube"  (containing  solid  calcium  chloride,  which  has 
the  property  of  absorbing  moisture).  Make  sure  that  the  apparatus  is  air-tight 
and  wrap  the  flask  in  a  towel. 

4.  Set  up  the  remainder  of  the  apparatus  according  to  the  directions  of  the 
instructor.     Dry  the  thick-walled  glass  test-tube  that  is  to  contain  the  copper 
oxide  by  heating  it  gently.     When  it  is  cool  weigh  it  carefully,  together  with 
any  portion  of  the  apparatus  that  may  come  in  contact  with  the  copper  oxide. 
Place  in  the  tube  about  1  gram  of  copper  oxide,  wipe  off  any  particles  that  are 
not  in  the  portion  of  the  tube  that  is  to  be  heated.     Weigh  again  carefully  to 
obtain  the  weight  of  copper  oxide  used.     Attach  the  apparatus  to  the  hydrogen 
generator,  pour  about  40  cc.  concentrated  hydrochloric  acid  down  the  thistle  tube, 
and  allow  the  hydrogen  to  pass  through  the  apparatus  until  it  has  swept  out 
the  air.      (Caution.     Do  not     place  a  flame  near  the  outlet  nor  heat  the  oxide 
while  the  apparatus  contains  a  mixture  of  oxygen  and  hydrogen.     A  dangerous 
explosion  might  result.)      Collect  the  gas  in  small  test  tubes  by  displacement 
of  water  and  ignite  it.    Explain  how  this  test  may  be  used  to  determine  when  the 
hydrogen  is  no  longer  mixed  with  oxygen. 

5.  When  pure  hydrogen  is  passing  over  the  copper  oxide,  begin  to  heat  the  oxide 
very  gently  with  a  small  flame  and  continue  to  heat  cautiously  until  all  the  oxide 
changes  color.    If  moisture  collects  in  the  farther  end  of  the  tube,  drive  it  out  by 
heating  the  tube  carefully.     Question.    Where  does  this  moisture  come  from? 

6.  Allow  the  tube  to  cool  in  the  current  of  hydrogen,  and  weigh  it.  If  you 
have  time,  check  this  result  at  once  by  repeating  the  heating  in  the  current  of 
hydrogen  and  then  weighing;   if  not,  set  the  tube  aside  in  order  that  you  may  do 
so  if  the  results  of  the  following  calculations  are  unsatisfactory. 

7.  Calculate  the  number  of   (a)  grams;    (b)   gram  atoms  of  copper  that  are 
combined  with  1  gram  atom  of  oxygen.    What,  then,  is  the  formula  of  this  oxide 
of  copper?    Repeat  the  experiment  if  your  results  are  inconclusive. 

8.  Write  the  equation  for  the  reduction  of  copper  oxide  by  hydrogen,  and 
interpret  in  terms  of  (a)  atoms  and  molecules;    (b)  gram  atoms  and  mols,  and 
(c)  grams. 

9.  Calculate  the  percents  of  copper  and  of  oxygen  in  this  oxide  of  copper 
(a)    from  your  experimental   results;     (b)    from  the   formula  and  the  atomic 
weights  of  copper  and  oxygen. 

10.  Problems.     ( i)  The  formulas  of  cuprous  oxide  and  cupric  oxide  are  Cu2O 
and  CuO,  respectively.     Write  equations  for  the  reactions  between  the  heated 
oxides  and  hydrogen  to  form  copper  and  steam.    What  weight  of  copper  would 
be  obtained  from  one  gram  of  each  oxide?     What  is  the  percentage  composition 
of  each  oxide? 

(2)  What  weight  of  water  could  be  obtained  from  1  gram  of  cupric  oxide? 
What  volume  would  this  amount  of  water  occupy  at  1  atmosphere  pressure  and 
(a)  4°  C;    (b)  273°  C?   What  volume  of  hydrogen  at  273°  C  and  1  atmosphere 
pressure  is  required  to  form  this  amount  of  water?     Is  this  the  amount  that 
would  be  used  in  an  experiment  similar  to  the  one  actually  performed? 

(3)  What  are  the  formulas  of  cuprous  sulfide,  cupric  sulfide,  and  hydrogen 
sulfide?     Cuprous  and  cupric  sulfides  are  also  reduced  to  copper  when  they  are 
heated  in  a  current  of  hydrogen;    hydrogen  sulfide  is  formed.     Write  equations 
for  the  reactions. 


14 


ASSIGNMENT  5 

THE  REACTION  BETWEEN  AN  ACID  AND  A  BASE  IN  SOLUTION 
CONCENTRATION  IN  TERMS  OF  MOLS  PER  LITER 

References.     Hildebrand,    Chapter   V,   pages   76-80,    Chapter   VIII,   pages    105 

and  106. 

1.  In  this  Assignment  we  shall  study  the  reaction  between  sodium  hydroxide 
and  hydrochloric  acid  in  solution  and  shall  determine  the  amount  of  salt  that  is 
formed  from  a  measured  volume  of  sodium  hydroxide  solution.    In  the  preceding 
Assignments   the   amount  of   a   substance  was   determined  by   weighing  or  by 
measuring  the  volume  of  the  pure  substance.     It  is  often  more  convenient  to 
determine  the  quantity  of  a  substance  by  measuring  the  volume  of  a  solution 
which  contains  a  known  amount  of  the  substance  in  a  unit  volume  of  the  solution. 
The    amount    of    the    substance    in    a    unit    volume   of    solution    is    called    the 
concentration.     Question.     If  the  concentration  of  a  salt  solution  is  known  to  be 
10  g.  per  liter,  what  volume  would  you  measure  out  in  order  to  have  0.5  g.  of  salt? 

2.  It  is  necessary  first  to  examine  separately  the  properties  of  the  three  solutions, 
the    base,    acid    and    salt.      Experiment.      Prepare    dilute    solutions    of    sodium 
hydroxide,    NaOH,   and  of   hydrochloric   acid,   HC1,   by   diluting  5   cc.   of   the 
laboratory  solution  of  each  with  50  cc.  of  distilled  water,  and  also  a  dilute  solution 
of  NaCl  by  dissolving  between  one  and  two  grams  of  the  salt  in  50  cc.  of  distilled 
water. 

To  10  cc.  portions  of  each  of  the  three  solutions  add  a  few  drops  of  litmus 
solution.  Repeat  using  phenolphthalein. 

Taste  each  solution  by  dipping  a  glass  rod  into  the  liquid  and  touching  it  to 
the  tongue.  (Caution.  Do  not  taste  any  substance  in  the  laboratory  unless 
directed  to  do  so.) 

Test  a  drop  of  each  solution  in  a  colorless  gas  flame  by  means  of  an  iron  (or 
platinum)  wire.  A  yellow  flame  proves  the  presence  of  sodium. 

Evaporate  to  dryness  in  a  casserole  1  cc.  of  HC1  solution,  and  of  NaCl  solution. 
In  each  case  examine  if  there  is  a  residue.  Question.  What  conclusion  can  you 
draw  in  regard  to  the  volatility  of  water  and  hydrogen  chloride  as  compared  to 
sodium  chloride?  What  result  would  you  predict  if  a  solution  containing  both 
NaCl  and  HC1  were  evaporated?  Pure  sodium  hydroxide  is  a  stable  non-volatile 
substance  which  would  be  left  as  a  solid  when  a  solution  containing  it  is 
evaporated  to  dryness.  (Caution.  Do  not  evaporate  alkaline  solutions  to  dryness. 
Glass  and  porcelain  are  slowly  attacked  by  hot  concentrated  alkaline  and  a  porcelain 
dish  is  spoiled  if  an  alkali  residue  is  heated  strongly  in  it.) 

Summarize  in  a  table  the  properties  of  the  three  solutions  examined  above. 

3.  Experiment.     Take  40  cc.  of  your  laboratory  NaOH  solution,  place  it  in  a 
clean  half  liter  flask  and  dilute  with  440  cc.  of  distilled  water.    Shake  the  flask  in 
order  that  the  solution  shall  be  uniform  throughout.     Cork  the  flask  and  label  it 
"NaOH  solution  for  Assignments  5,  6  and  7."     Question.    What  approximately 
is  the  ratio  of  the  initial  volume  of  the  sodium  hydroxide  solution  to  the  final 
volume? 

4.  Dry  a  porcelain  dish  and  watch-glass  large  enough  to  cover  it.    Weigh  the 
dish  and  watch-glass  to   10  mg.     Measure  out  in  your  graduated  cylinder  as 
accurately  as  possible  50  cc.  of  the  sodium  hydroxide  solution  prepared  above. 
Pour  the  solution  into  the  weighed  evaporating  dish,  and  sufficient  phenolph- 
thalein to   give  a  pink   color  and  then  add   small  portions  of  your  laboratory 
hydrochloric   acid,   stirring  after   each   addition,   until  the   solution   is   colorless. 
Approximately  5  cc.  of  the  acid  will  be  required  to  give  the  colorless  solution. 
If  the  addition  of  the  last  portion  of  acid  is  made,  a  few  drops  at  a  time,  it  will 

f  151 


be  observed  that  the  color  changes  abruptly.     An  excess  of  1  cc.  of  HC1  may 
now  be  added. 

5.  Place  the  dish  containing  the  solution  on  your  wire  gauze  and  heat  until 
the  solution  begins  to  boil.    Reduce  the  size  of  the  flame  and  allow  the  solution 
to  boil  gently,  or  to  evaporate   slowly  just  below  the  boiling  point,  until  the 
bottom  of  the  dish  is  covered  with  solid  material.     Then,  to  avoid  loss  from 
bumping  and  spattering  during  the  evaporation  to  dryness,  cover  the  dish  loosely 
with  the  watch-glass,  leaving  an  open  space  at  one  side,  and  continue  the  heating, 
first  with  a  small  flame,  then  more  strongly  until  no  further  trace  of  water-vapor 
is  expelled.     Allow  the  dish  and  residue  to  cool  for  10  minutes  while  covered 
with  the  watch-glass  and  again  weigh  to  10  mg.    Heat  the  dish  and  residue  gently 
for  five  minutes  longer,  let  cool,  and  weigh  again.    If  the  two  weights  are  not  the 
same  within  20  mg.   repeat  this  process  until  two  weights  are  obtained  which 
check  to  20  mg. 

6.  Dissolve  a  portion  of  the  residue  in  a  small  amount  of  water  and  test  as 
in  Paragraph  2.     State  what  evidence  you  have  that  a  reaction  has  taken  place 
between  the  acid  and  base,  and  that  the  solid  residue  obtained  is  sodium  chloride. 
Write  the  equation  for  the  neutralisation  reaction  between  NaOH  and  HC1  and 
interpret  it  in  terms  of   (a)   molecules,   (b)  mols,   (c)   grams  of  the  substances 
involved. 

7.  Calculations.     From  the  weight  of  NaCl  found  in  Paragraph  5,  calculate 
(a)  the  weight  of  NaOH  which  must  have  been  present  in  the  50  cc.  of  NaOH 
solution,   (b)  the  number  of  grams  of  NaOH  in  1  cc.  of  the  solution,   (c)  the 
concentration  of  the  NaOH  in  grams  per  liter?     As  we  shall  see  in  the  next 
assignment  this  value  may  not  be  very  accurate.     In  addition  to  errors  such  as 
that  made  in  measuring  out  the  volume  of  the  NaOH  solution,  the  laboratory 
NaOH  contains  small  amounts  of  impurities,  as  NaCl. 

8.  The  reaction  just  considered  is  typical  of  the  reaction  between  any  acid  and 
any  base.    In  every  case  H  of  the  acid  unites  with  OH  of  the  base  to  form  water. 
Write  the  equations  for  the  reactions  between  the  following  bases  and  acids,  and 
interpret  each  equation  in  terms  of  mols  of  the  substances  involved : 

Sodium  hydroxide  and  nitric  acid 

Sodium  hydroxide  and  sulfuric  acid  (to  form  two  molecules  of  water) 

Barium  hydroxide  and  hydrochloric  acid 

Barium  hydroxide  and  sulfuric  acid. 

9.  Since  the  mol  is  a  convenient  unit  of  weight  to  use  in  studying  chemical 
reactions,  concentration  is  often  expressed  in  terms  of  the  number  of  mols  of 
substances  in  a  liter  of  solution.    {Definition.     A  solution  which  contains  in  one 
liter   one   mol   of   dissolved   substance   is   called   a   molal   solution;    one   which 
contains  in  one  liter  one-tenth  mol  of  dissolved  substance  is  called  a  tenth  molal 
solution,  etc.     Molal  hydrochloric  acid  is  designated  thus:   M  HC1;   tenth  molal 
sulfuric  acid  would  be  written  0.1   M  H2SO4,  etc.     Questions.     Calculate  the 
number  of  mols  of  NaCl  obtained  in  Paragraph  5.     How  many  mols  of  NaOH 
were  there  in  50  cc.  of  solution  ?    What  then  is  the  concentration  of  your  NaOH 
in  mols  per  liter?     If  the  laboratory  solution  were  exactly  6  M  (which  probably 
is  not  the  case)  what  would  this  concentration  be,  as  a  result  of  the  twelve  fold 
dilution  in  Paragraph  3? 

10.  Problems,     (i)   How  many   (a)   mols,   (b)  grams  of  sulfuric  acid  are  in 
50  cc.  of  0.2  M  H2SO4? 

(2)   What  is  the  concentration  in  mols  per  liter  of  a  solution  which  contains 
5.8  g  of  NaCl  in  125  cc.  of  solution. 


16 


ASSIGNMENT  6 

TITRATION    OF    SOLUTIONS   OF   ACIDS    AND    BASES  :       AN    ILLUSTRATION 
OF  VOLUMETIC  ANALYSIS 


Reference.     Hildebrand,  Chapter  VIII,  pages  106-107. 

1.  In  Assignment  5  we  learned  that  an  acid  and  a  base  in  solution  will  neutralize 
each  other.     In  Assignment  6  we  shall  see  how  this  reaction  can  be  used  in 
determining  the  concentration  of  one  of  these  solutions  when  that  of  the  other  is 
known.     The  operation  is  called  a  titration.     It  is  evident  that  a  pure  solution 
of  a  salt  may  be  prepared  by  mixing  the  corresponding  acid  and  base  in  exactly 
the  right  proportion;    this  end-point  in  the  titration  is  determined  by  means  of  a 
suitable  indicator     .The  relative  concentration  of  the  two  solutions  can  be  calcu- 
lated when  the  relative  volumes  of  the  two  solutions  used  in  the  titration  are 
known.      Questions.     How   many   mols   of   NaOH   are   required   to   neutralize 
exactly  0.01  mol  of   (a)   HC1,   (b)   H,SO4?     How  many  cc.  of  0.50  M  NaOH 
solution  will  exactly  neutralize  0.01  mol  of  (a)  HC1,  (b)  H2SO,.? 

2.  Experiment.     Prepare  300  cc.  approximately  0.5  M  HC1  from  your  labora- 
tory 6  M  solution,  place  it  in  a  flask  and  shake  it.     Cork  the  flask  and  label  it 
"HC1   solution   for  Assignments   6  and   7."     Clean   a   small   flask   and   label   it 
"known  H2SO4  solution/'  rinse  the  flask  with  distilled  water  and  set  it  aside  to 
drain  in  order  to  have  it  ready  for  use  later  in  this  Assignment. 

3.  Experiment  on   the  determination  of  an  end-point  and  the  choice  of  the 
indicator  to  be  used  in  the  titration.    Dissolve  approximately  0.5  g.  NaCl  in  about 
50  cc.  water,  add  2  drops  phenolphthalein  and  stir  the  solution.     Add  NaOH 
solution  (approximately  0.5  M)  drop  by  drop,  stirring  after  each  drop  is  added, 
and  note  how  many  drops  are  needed  to  give  a  distinct  color.     Then  determine 
how  many  drops  of  your  0.5  M  HC1  solution  are   required  to  decolorize  the 
solution.     Note.     Small  drops  are  conveniently  added  from  a  glass  tube  drawn 
out  to  a  point,  the  solution  being  held  in  the  tube  by  placing  the  finger  over  the 
upper  end  of  the  tube.    Be  sure  that  the  tube  is  clean ;   before  using,  rinse  it  once 
with  the  solution.    Repeat  the  experiment  with  litmus  solution  instead  of  phenol- 
phthalein, and  determine  whether  the  change  in  color  gives  a  satisfactory  end- 
point   for   the   titration   of   HC1   and   NaOH   solutions.     Finally   determine  the 
nature  of  the   end-point  with  each  indicator  when  about   1   g.   sodium  sulfate, 
Na2SO4-  10  H2O  is  used  instead  of  NaCl. 

4.  The  volumes  of  the  solutions  used  in  a  titration  are  measured  by  means  of 
burettes;  these  are  uniform  glass  tubes  graduated  in  cubic  centimeters  and  tenths 
(or  fifths)  of  cubic  centimeters.    Two  burettes  will  be  placed  on  your  desk  before 
the  beginning  of  the  Assignment.     At  the  end  of  the  period  rinse  them  with 
distilled  water  and  leave  them  on  your  desk. 

5.  Experiment.     Fill  each  burette  with  distilled  water.     Air  may  be  removed 
from  the  small  tube  below  the  pinch-cock  by  tilting  the  tip  upward  and  allowing 
the  liquid  to  flow  through  the  pinch-cock.     Question.     Why  is  it  necessary  to 
remove  any  bubble  of  air  trapped  in  this  small  tube?     Practice  reading  a  burette: 
bring  your  eye  to  the  same  level  as  the  liquid  and  note  the  reading  of  the  burette 
corresponding  to  the  bottom  of  the  meniscus;    repeat  until  consecutive  readings 
check  to  better  than  0.05  cc.     Question.    Why  is  it  important  to  have  the  eye  at 
the  same  level  as  the  liquid  before  making  a  reading?     Never  attempt  to  adjust 
the  volume  of  the  solution  in  a  burette  so  that  the  reading  will  be  some  exact 
amount.    Allow  the  water  to  flow  slowly  out  of  the  burette.    If  drops  remain  on 
the  inner  surface  of  a  burette,  exchange  it  at  the  office  for  a  clean  one. 

6.  Rinse  one  burette  with  a  little  of  your  approximately  0.5  molal  HC1  solution, 
and  fill  the  burette  with  this  solution.     Rinse  and  fill  the  other  burette  with  the 
0.5  molal  NaOH  solution. 

[  171 


7.  Record  the  readings  of  the  burettes  side  by  side  in  your  note-book ;    run 
about  15  cc.  of  the  acid  solution  into  a  clean  beaker  or  flask  standing  on  white 
paper,  and  record  the  final  burette  reading  under  the  initial  reading.     Add  two 
drops  of  phenolphthalein,  and  about  20  cc.  distilled  water.     Then   run  in  the 
sodium  hydroxide  solution  from  the  other  burette,  a  little  at  a  time,  and  towards 
the  end,  very  carefully,  a  drop  or  two  at  a  time,  stirring  the  mixture  constantly, 
until  the  faintest  perceptible  permanent  pink  color  is  obtained.     Wash  down  the 
inside  of  the  beaker  by  means  of  a  jet  of  water  from  the  wash  bottle.     If  two 
much  of  the  basic  solution  is  added,  decolorize  the  solution  by  adding  a  little  of 
the  acid  and  determine  the  end-point  again.     Record  the  final  readings  of  each 
burette,  and  the  actual  volumes  of  each  solution  used  in  titration.     Calculate  the 
volume  of  sodium  hydroxide  necessary  to  neutralize  one  cubic  centimeter  of  the 
hydrochloric  acid. 

8.  Repeat  this  experiment,  using  about  20  cc.  of  the  acid  solution,  and  in  each 
case  make  the  same  calculations.    Do  not  fill  up  the  burette  each  time  unless  there 
is  not  enough  solution  in  it  for  the  titration. 

9.  Questions.    If  the  error  in  measuring  out  a  volume  of  solution  by  means  of  a 
burette  is  0.10  cc.  what  is  the  percentage  error  if  1  cc.  of  solution  is  measured? 
If  20  cc.  of  solution  are  measured?     Why  is  it  important  not  to  use  less  than 
10  cc.  in  any  titration? 

10.  Compare  the  volume  ratios  calculated  from  two  titrations.     If  the  result 
differs  from  the  average  by  more  than  1%,  perform  additional  titrations  until 
you  are  satisfied  that  you  have  determined  the  volume  ratio  with  an  accuracy 
better  than  1%. 

11.  Questions,     (a)  From  an  examination  of  the  equation  for  the  neutraliza- 
tion of  sodium  hydroxide  by  hydrochloric  acid  state  the  ratio  of  the  number  of 
mols  of  acid  and  base  added  to  the  beaker  when  exact  neutrality  was  reached. 

(b)   From  your  average  volume  ratio  state  which  solution,  acid  or  base,  is  the 
more  concentrated,  and  what  is  the  ratio  of  the  concentrations. 

12.  Take  your  clean,  dry,  labelled  flask,  Paragraph  2,  to  the  office  to  obtain  a 
sulfuric  acid  solution  of  known  concentration. 

13.  Empty  the  HC1  out  of  the  burette,  rinse  it  with  about  5  cc.  of  the  sulfuric 
acid    solution    of    known    concentration,    and    fill    it    with    the    sulfuric    acid 
solution.     Determine  the  volume  ratio  as  before  from  three  (or  more)  titrations. 
From  the  volume  ratio  and  the  reaction  between  sodium  hydroxide  and  sulfuric 
acid  calculate  the  concentration  of  the  sodium  hydroxide  solution  in  mols  per 
liter.     Calculate  also  the  concentration  of  your  hydrochloric  acid  solution. 

14.  Make  a  list  of  the  sources  of  error.     (Many  of  them  have  been  mentioned 
in  the  above  directions.) 

15.  Save  the  remainder  of  the  NaOH  and  HC1  solutions,  whose  concentration 
you  have  determined,  in  corked  flasks  for  use  in  Assignment  7. 

16.  Problems,     (i)   How  many  cc.  of  0.01  M  Ba(OH)2  will  be  required  to 
neutralize  10  cc.  of  0.5  M  HC1? 

(2)   What  is  the  concentration  in  mols  per  liter  of  a  sulfuric  acid  solution, 
25  cc.  of  which  neutralizes  20  cc.  of  0.20  M  NaOH? 


[18 


ASSIGNMENT  7 

VOLUMETIC  ANALYSIS,  CONTINUED:  THE  DETERMINATION  OF  THE 
EQUIVALENT  WEIGHT  OF  AN  UNKNOWN  ACID 


Reference.    Hildebrand,  Chapter  VIII,  pages  108-111. 

1.  In   Assignment  7  there   will   be  introduced   another   unit  of   quantity,   the 
equivalent.     This  unit  and  the  corresponding  unit  of  concentration,  equivalents 
per  liter,  are  frequently  more  convenient  than  the  units,  mol  and  mols  per  liter 
The  mol  and  equivalent  are  identical  for  HC1,  HNO3,  NaOH,  NaCl,  etc.,  but  a 
mol  of  H2SO4,  Ba(OH)2  or  Na2SO4,  etc.,  contains  two  equivalents.    A  solution 
which  contains  one  equivalent  in  a  liter  is  called  a  normal  solution  and  is  desig- 
nated i  N.    The  convenience  of  this  unit  of  concentration  depends  upon  the  fact 
that  when  one  equivalent  of  any  acid   reacts  with  one  equivalent  of  any  base 
the    resulting    solution    contains    one    equivalent    of    the    corresponding    salt. 
Question.     What  is  the  normal   concentration  of  a  molal  solution   of  H2SO4, 
NaOH    and    Na2SO4    respectively?      How    many    equivalents    of    acid    can    be 
neutralized  by  10  cc.  of  0.1  N  NaOH?     Calculate  the  normal  concentrations  of 
your  solutions  of  NaOH,  HC1  and  H2SO4  used  in  Assignment  6. 

2.  Weighed  .portions  of  an  unknown  solid  acid  will  be  titrated  with  your  NaOH 
solution.     By  means  of  the  concentration  of  the  NaOH  solution  (determined  in 
Assignment  6)  the  number  of  grams  in  one  equivalent  of  the  acid  will  be  calcu- 
lated.   The  correctness  of  this  result  depends  of  course,  on  the  accuracy  of  your 
work  in  Assignment  6.     It  is  to  be  noted  that  if  you  had  started  with  a  solid  and 
of  known  equivalent  weight,  you  could  have  used  the  results  obtained  in  this 
Assignment  to  determine  the  concentration  of  the  sodium  hydroxide  solution. 

3.  Experiment.     Obtain  from  the  office  a  sample  bottle  containing  approxi- 
mately 2  grams  of  a  crystalline  acid  of  unknown  composition.     Clean  two  %  liter 
flasks  and  label  them  No.  1  and  No.  2.    Weigh  the  sample  bottle  with  its  cork  and 
contents.     Remove  the  cork,  taking  care  that  none  of  the  solid  which  may  be 
sticking  to  it  drops  off,  and  shake  about  one  gram  into  flask  No.  1.     Replace  the 
cork  and  weigh  again.     Shake  the  remainder  of  the  sample  into  flask  No.  2,  and 
weigh  the  empty  sample  bottle  and -cork.    All  weighings  should  be  made  to  5  mg. 

4.  Dissolve  the  contents  of  each  flask  in  about  50  cc.  of  distilled  water,  add 
two  drops  of  phenolphthalein  and  titrate  to  the  appearance  of  the  first  pink  color 
with  the  sodium  hydroxide  solution  prepared  in  Assignment  8.     If  you  should 
pass  the  end-point  in  a  titration  add  from  a  second  burette  an  acid  solution  of 
known  concentration  until  the  pink  color  is  discharged,  and  again  titrate  to  the 
end-point  with  the  sodium  hydroxide. 

5.  For   each   sample   calculate   the   number   of    cubic   centimeters   of   NaOH 
solution  needed  to  neutralize  1  gram  of  the  acid.     If  the  results  differ  by  more 
than  2%  repeat  the  experiment. 

6.  From  your  two   (or  more)   measurements  obtain  an  average  value  of  the 
number  of  cubic  centimeters  of  sodium  hydroxide  solution  per  gram  of  acid. 
From  this  average  value  and  the  concentration  of  the  NaOH  solution  calculate 
the  number  of  mols  of  NaOH  needed  to  neutralize  1  gram  of  the  acid. 

7.  Questions.     How  many  equivalents  are  in  one  gram  of  the  acid?     How 
many  grams  are  in  one  equivalent  of  the  acid? 

Report  this  value  to  your  instructor  at  once.  If  your  result  is  unsatisfactory 
check  your  calculations  in  Assignments  6  and  7  and  repeat  as  much  of  the  work 
as  is  necessary. 

8.  Problems,     (i)   Among  the  acids  suitable  for  this  Assignment  are:    oxalic 
H2C2O4  •  2H2O  ;  citric,  H3C6H5O7  •  H2O  ;  tartaric,  HX4H4Oe  and  potassium  acid 
sulfate,  KHSO4.     Write  the  equations  for  the  reaction  between  NaOH  and   (a) 

[19] 


oxalic  acid  to   form  Na2C2O4,  and    (b)    KHSO4;  and   calculate   the  equivalent 
weight  of  the  acid  in  each  of  these  reactions. 

(2)  What  is  the  normal  concentration  of  a  sulfuric  acid  solution,  25  cc.  of 
which  neutralizes  20  cc.  of  0.20  AT  NaOH? 

(j)  Chemically  pure  ("C.  P.")  sulfuric  acid,  nitric  acid,  hydrochloric  acid 
and  ammonia,  as  supplied  by  the  manufacturers,  are  concentrated  aqueous 
solutions  of  these  substances.  The  concentration  of  each  solution  is  guaranteed 
not  to  be  less  than  a  certain  minimum  value,  and  this  is  tested  by  measuring  the 
density  (or  the  specific  gravity).  The  following  table  contains  the  density  and 
the  percentage  composition  by  weight  of  the  concentrated  laboratory  reagents. 

Density  Concentration, 

g.  per  cc.  %  by  weight  Equivalents  per  liter 

H2SO4  1.84  95.6%  H2SO4 

HNO3  1.42  69.8%  HNO3 

HC1    '  1.19  37.2%  HC1 

NH4OH  0.90  57.5%  NH4OH 

For  each  solution  calculate  the  normal  concentration  and  record  the  results  in 
the  fourth  column  of  the  table. 

Caution!  The  concentrated  acids,  especially  sulfuric  and  nitric,  produce 
dangerous  burns  and  should  not  be  used  carelessly.  When  you  are  directed  to 
use  one  of  these  acids  carefully  pour  just  enough  of  it  for  your  experiment  into 
a  dry  beaker. 


20 


SECTION  II 

IONIC  THEORY 

RAPID  REVERSIBLE  REACTIONS  AND  EQUILIBRIUM 


ASSIGNMENT  21 
SOLUTIONS  OF  STRONG  ELECTROLYTES.     IONIC  EQUATIONS 

Reference.     Hildebrand,  Chapter  X,  pages  124-137. 

1.  This  Assignment,  which  contains  no  experimental  work,  is  introduced  in 
order  that  the  student  may  become  familiar  with  the  fundamental  ideas  under- 
lying the  Ionic  Theory  before  proceeding  to  use  these  ideas  in  the  following 
Assignments. 

2.  The  use  of  the  terms  acids  and  bases  in  designating  two  distinct  groups  of 
substances  implies  that  the  members  of  each  group  have  a  set  of  properties  in 
common.    The  properties  characteristic  of  all  acid  solutions,  which  were  observed 
in  Assignment  5,  are  ascribed  to  a  substance  called  hydrogen  ion,  represented 
by  the  symbol  H+;   and  those  of  basic  solutions  to  the  substance  hydroxide  ion, 
written  OH~.     In  addition  to  the  properties  of  the  hydrogen  ion,  each  acid  in 
solution  has  a  group  of  properties  different  from  those  of  any  other  acid  but 
common  to  solutions  of  all  salts  of  that  acid.    Thus,  hydrochloric  acid  has  a  set 
of  properties  which  is  characteristic  of  solutions  of  all  chlorides  and  is  ascribed 
to  a  substance  called  chloride  ion,  Cl~.    Likewise,  a  solution  of  sodium  hydroxide 
has  a  group  of  properties  which  is  characteristic  of  solutions  of  all  sodium  salts 
and  which  is  attributed  to  the  sodium  ion,  Na+.    Questions.    What  is  the  evidence 
from  freezing  point  data  that  there  are  approximately  2  mols  of  substance  pres- 
ent when  one  mol  of  hydrogen  chloride  is  dissolved  in  water?     How  does  the 
electrical  conductivity  of  hydrochloric  acid  solution  support  the  idea  that  the 
molecule  of  hydrogen  chloride  in  solution  is  broken  up  into  two  new  substances? 
In  what  way  does  a  chloride  ion  differ  from  an  atom  of  chlorine,  and  a  hydrogen 
ion  differ  from  an  atom  of  hydrogen?     List  briefly  differences  in  properties  of 
the  substances  hydrogen  ion  and  hydrogen  gas. 

3.  Many  substances  in  dilute  solution  may  be  considered  as  completely  ionized. 
These  substances  are  called  strong  electrolytes  and  include : 

Practically  all  salts 

A  few  acids  as  HC1,  HNO3  and  H2SO4,  and 

A  few  bases  as  NaOH,  KOH  and  Ba(OH)2. 

The  student  should  memorize  this  list  and  should  form  the  habit  of  thinking 
of  solutions  of  strong  electrolytes  in  terms  of  the  ions  present  and  not  merely  in 
terms  of  the  specific  solid,  liquid  or  gas  used  in  making  the  solution.  It  is  impor- 
tant to  realize,  however  that  strong  electrolytes  are  not  ionized  in  the  solid  or 
gaseous  state ;  thus,  while  hydrochloric  acid  and  sodium  chloride  solutions  consist 
of  H+  and  Cl~  and  Na+  and  Cl~,  respectively,  gaseous  HC1  and  solid  NaCl  are 
not  ionized,  and  each  has  its  own  specific  properties.  Questions.  What  are  the 
principal  substances  present  in  each  of  the  following  solutions  and  what  is  the 
approximate  concentration  of  each  substance  in  mols  per  liter : 

(/)  A  solution  which  contains  0.1  mol  of  H2SO4  in  1  liter.  (Answer.  H+  and 
SO4~"  at  concentrations  0.2  M  and  0.1  M,  respectively). 

(-2)  A  solution  which  contains  0.2  mol  of  NaOH  in  1  liter. 

(5)  A  solution  which  contains  0.1  mol  of  Na0SO4  and  0.1  mol  of  NaCl  in 
1  liter. 

[21] 


(4)   A  solution  which  is  made  by  mixing  equal  volumes  of  (i)  and  (2). 

4.  Ionic  equations.  Having  realized  what  substances  are  present  in  solutions 
of  strong  electrolytes  we  are  now  in  a  position  to  interpret  reactions  involving 
such  solutions.  We  shall  first  consider  what  is  the  ionic  reaction  when  a  strong 
acid  neutralizes  a  strong  base,  and  shall  take  as  an  example  the  reaction  between 
sulfuric  acid  and  sodium  hydroxide  solutions.  The  equation 
H,SO4  +  2NaOH  =  2H2O  +  Na2SO4 

is  a  statement  of  the  reaction  that  takes  place  whenever  H2SO4  and  NaOH 
solutions  are  mixed,  and  we  have  seen  that  it  may  be  interpreted  in  terms  of 
molecules,  and  of  mols,  equivalents,  grams,  or  any  other  weight  units.  It  records 
no.  experimental  details,  such  as  volumes  or  concentrations  of  solutions,  or  the 
use  of  either  reagent  in  excess;  and  additional  notes  are  necessary  when  such  a 
record  is  desired. 

This  equation  can  be  interpreted  according  to  the  ionic  theory  (a)  by  adding 
a  note  that  H2SO4,  NaOH  and  Na2SO4  are  strong  electrolytes  and  that  H2O  is  a 
weak  electrolyte,  or  (b)  by  rewriting  the  equation  to  give  the  same  information: 

(2H+  +  SO4~)  +  (2Na+  +  2OH~)  =  2H2O  +  2Na+  +  SO4~. 
It  is  evident  that  the  Na+  and  SO4~~  shown  to  be  in  the  final  solution  were  present 
in  the  two  initial  solutions,  and  that  these  substances  have  undergone  no  change 
during  the  reaction.  It  is  incorrect  to  say  that  "sodium  ion  and  sulfate  ion  have 
combined/'  The  reaction  that  has  taken  place  is  simply  the  formation  of  the  weak 
electrolyte  water: 


5.  The  statement  that  a  substance  is  a  weak  electrolyte  is  equivalent  to  saying 
that  the  ions  of  that  substance  cannot  exist  in  the  presence  of  each  other  except 
at  low   concentrations   and  it  is  obvious   that  any   reaction  which  involves   the 
formation  of  a  weak  electrolyte  from  its  ions  may  be  expected  to  take  place. 
Accordingly,  we  shall  be  able  to  predict  a  number  of  reactions  when  we  have 
classified  the  substances  involved  as  strong  or  weak  electrolytes  in  solution. 

6.  Another   reaction   which   involves   the  ions   of  a   strong  electrolyte   is   the 
precipitation  of  a  sparingly  soluble  salt.    Again  the  statement  that  a  salt,  such  as 
silver  chloride,  AgCl  ;  is  sparingly  soluble  is  equivalent  to  saying  that  its  ions 
cannot  exist  together  in  solution  except  at  the  low  concentrations  corresponding 
to  the  solubility  of  the  salt.     Write  the  ionic  equation  for  the  reaction  that  takes 
place  when  dilute  solutions  containing  equivalent  amounts  of  NaCl  and  AgNO3 
are  mixed.     What  substances  are  present  at  high  concentrations  in  each  of  the 
initial   solutions   and   in   the   final   solution?     Write  the   ionic  equation   for  the 
reaction.     Consider  next  the  case  in  which  silver  nitrate  solution  in  excess  is 
added  to  a  sodium  chloride  solution,  answer  the  same  question  and  write  the 
ionic  equation.     It  should  be  obvious  that  the  reaction  is  the  same  in  both  cases, 
and  that,  therefore,  the  two  equations  should  be  identical. 

7.  Does  anything  happen  in  the  following  experiments  ?     A  dilute  solution  of 
sodium    chloride    is    mixed    with    a    dilute    solution    of    (i)    potassium    nitrate, 
(2)  nitric  acid? 

8.  Problems,     (i)   Write  ionic  equations  for  the  following  reactions:    (a)    A 
precipitate  of  barium  sulfate,  BaSO4,  is  formed  by  mixing  solutions  of  barium 
chloride  and  sulfuric  acid,     (b)  A  sodium  sulfate  solution  is  evaporated  to  dry- 
ness.     (c)   Hydrogen  chloride  gas  is  dissolved  in  water. 

(2)  The  solubility  of  lead  iodide,  PbI2,  is  0.002  mol  per  liter  at  18°.  What 
is  the  concentration  of  the  ions  present  in  a  saturated  solution? 

(j)  If  the  freezing  point  of  water  is  lowered  approximately  1.86°  per  mol  of 
substance  in  solution  in  1000  grams  of  water,  what  is  the  freezing  point  of  a  0.1 
molal  solution  of  (a-)  a  non-electrolyte,  (b)  hydrochloric  acid,  (c)  barium 
chloride,  BaQ2? 

[221 


ASSIGNMENT  22 

STRONG  AND  WEAK  ACIDS.    THE  USE  OF  INDICATORS  TO  MEASURE 
HYDROGEN  ION  CONCENTRATION 


Reference.     Hildebrand,  Chapter  X,  pages  138-142. 

1.  Indicators  can  be  used  to  measure  the  concentration  of  hydrogen  ion,  or  of 
hydroxide  ion,  in  a  solution.     The  color  change  for  each  indicator  occurs  in  a 
definite  range  of  concentrations  of  hydrogen  ion   (or  hydroxide  ion),  which  is 
characteristic  of  the  indicator.     In  the  present  Assignment,  by  using  solutions  of 
known  concentration  of  the  strong  acids  listed  in  Assignment  21,  we  shall  develop 
a  method  of  determining  approximately  the   concentration  of  hydrogen  ion  in 
any  acid  solution.    This  indicator  method  will  then  be  used  to  measure  the  concen- 
tration  of   hydrogen   ion   in   solutions   of  a   typical  weak   acid.     Two   reactions 
involving  this  acid  will  be  studied. 

2.  Experiment.     Prepare  60  cc.  N  HC1  by  adding  distilled  water  to  the  proper 
volume   of  the  6  N  laboratory   reagent  and   shaking  the   mixture.     From  this 
solution,  or  from  your  known  HC1  solution,  Assignment  6,  prepare  between  50 
and  100  cc.  of  0.10  N  HC1,  and  from  this  solution  prepare  50  to  100  cc.  of  0.001 
N  HC1.    On  account  of  the  error  in  measuring  volumes  by  means  of  a  graduated 
cylinder,    and    the   probable   variation    of   the    concentration   of   the    laboratory 
solution    from    6    N,    the    concentrations    of    these    solutions    are    known    only 
approximately. 

3.  Pour  into  marked  test  tubes  10  cc.  of  each  HC1  solution  (N,  0.1  N,  0.01  N, 
0.001  N)  and  pour  10  cc.  of  water  into  a  fifth  test-tube.     Add  to  each  solution 
from  a  glass  tube  a  single  drop  of  methyl  violet  solution.  Hold  the  tubes  in  a  verti- 
cal position  over  a  piece  of  filter  paper,  look  down  through  the  surface,  record  the 
color  of  each  solution  and  note  the  smallest  concentration  of  hydrochloric  acid 
that   shows   with   this   indicator  a   color   different   from  that  of   water.     If   the 
indicator  solution  is  so  dilute  that  one  drop  does  not  produce  a  distinct  color  add 
1  or  2  more  drops,  but  record  the  number  of  drops,  and  use  the  same  number  in 
all  the  tubes.     State  how  the  indicator,  methyl  violet,  may  be  used  to  determine 
the  approximate  concentration  of  a  hydrochloric  acid  solution.     (The  color  in  the 
more  concentrated  solution  will  fade  on  standing.     It  may  be  restored  by  adding 
another  drop  of  the  indicator.)   Note.     Set  aside  the  remainder  of  the  0.01  N 
and  0.001  TV"  HC1  for  use  later  in  this  Assignment,  Paragraphs  5  and  8. 

4.  Experiment.     Repeat  the  experiment  with  nitric  acid  or  with  sulfuric  acid, 
using  the  same  concentrations  as  before  (N,  0.10  N,  0.01  N,  0.001  N).    Compare 
the  colors  obtained  with  the  different  acids.     Questions.    Are  the  colors  charac- 
teristic   for    each    acid?      If    not,    what    substance    determines    the    color?      If 
hydrochloric  acid  is  completely  ionized  what  conclusion  can  you  draw  with  respect 
to  the  ionization  of  nitric  acid  and  sulfuric  acid?     Note.     Save  the  0.001   N 
solution  for  later  use,  Paragraph  5. 

5.  If  you   have  performed  the  above   experiments   correctly  and  understood 
them  you  will  realize  that  the  indicator  methyl  violet  can  be  used  to  determine, 
approximately,     concentrations    of    hydrogen    ion    between    N    and    0.001     N. 
However,  while  the  concentration  of  H+  in  the  dilute  solution,  0.001  N,  (which  is 
often  written  10~3  N),  is  small  compared  with  the  normal  solution,  it  is  10,000 
times   as   great   as   in  pure   water.     We   shall  now   make  use   of   the   indicator 
methyl  orange  to  examine  solutions  in  which  the  concentration  of  H+  is  between 
10~3  N  and  that  of  pure  water,  10~7  N.     Experiment.     By  a  10  fold  dilution  of 
your  10-3  A7  HC1  solution,  prepare  50  or  100  cc.   10~4  N  HC1;  and  from  this 
prepare  a  10~5  N  solution.     Test  10  cc.  portions  of  these  three  solutions  and  of 
water  with  1  drop  of  methyl  orange.    Repeat  the  experiment  with  nitric  or  sulfuric 

[23] 


acid,  starting  with  your  10~3  A7"  solution.    Note.     Save  one  of  the  lO'1  N  solutions 
for  later  use,  Paragraph  8. 

6.  Summarize  your  results  with  the  two  indicators  in  a  table  which  shows  the 
color  obtained  at  various  concentrations  of  H+.  Show  your  table  to  the 
instructor. 

7 '.  The  weak  acid  which  wre  shall  now  proceed  to  study  is  acetic  acid, 
HC2H3O2.  It  is  a  white,  crystalline,  rather  volatile  solid  which  melts  near  room 
temperature  and  is  very  soluble  in  water.  The  formula  of  its  sodium  salt,  sodium 
acetate,  is  NaC2H3O2.  We  shall  abbreviate  these  formulas  to  HAc  and  NaAc, 
respectively.  Questions.  How  many  mols  of  acetic  acid  are  contained  in  1  liter 
of  6  N  acid?  How  many  mols  of  sodium  acetate  could  be  prepared  from 
this  quantity  of  acetic  acid? 

8.  Experiment.     From  the  laboratory  6  N  HAc  prepare  solutions  which  are 
approximately  N,  0.5  N,  0.05  N  and  0.01  N.     Place  a  10  cc.  portion  of  each  of 
the  acetic  acid  solutions  (N  to  .01  N)  in  a  labelled  test-tube,  place  10  cc.  distilled 
water  in  another  test-tube,  and  test  each  solution  with  methyl  violet  as  in  Para- 
graph   3.      For    comparison,    repeat    the   test    with    HC1    solutions    of    suitable 
concentrations.     (Note.     In  color  comparisons  it  is  not  safe  to  trust  the  memory, 
or  even  written  descriptions.)     Determine  the  lowest  concentration  of  acetic  acid 
at  which  the  color  with  methyl  violet  is  distinctly  different  from  that  with  water, 
and  estimate  approximately  the  concentration  of  hydrogen  ion  in  two  of  the  acetic 
acid  solutions,  say  in  the  N  and  O.I  N  solutions.     Repeat  the  experiment,  but 
use  methyl  orange  instead  of  methyl  violet,  and  estimate  the  concentration  of  H+ 
ion  in  the  O.I  N  solution  and  in  a  more  dilute  solution. 

9.  The  acetic  acid  in  the  solution  must  be  present  either  in  the  form  of  ions, 
H+  and  Ac~,  or  in  the  un-ionized  form,  HAc.    From  your  estimate  of  the  concen- 
tration of  H*  in  the  0.1  normal  solution,  calculate  the  fraction  of  the  acetic  acid 
which  is  ionized,  and  the  fraction  which  is  un-ionized.    The  fraction  of  the  acid 
which  is  in  the  form  of  ions  is  called  the  degree  of  ionization.     State  also  the 
concentrations  of  acetate  ion,  Ac~,  and  of  the  un-ionized  acid,  HAC,  in  the  0.1 
normal  acetic  acid  solution.  Is  acetic  acid  a  weak  or  a  strong  acid? 

10.  Calculations.     The  concentration  of  the  ions  in  acetic  acid  solutions  have 
been  determined  by  other  methods  more  accurately  than  is  possible  by  these  color 
experiments.     The  concentrations  of  hydrogen  ion  in  these  solutions  at  room 
temperature  are  given  in  the  following  table: 

Concentration  Concentration     Concentration     Concentration  of     Degree  of 
acid.  of  H+.  of  Ac~.  un-ionized  HAc.    Ionization. 

1  N  .004  N 

0.5  AT  .003  N 

r\N  .0013  N 

O.OIN  .0004  AT 

Fill  in  the  remaining  columns  of  the  table  and  show  your  table  to  the  instructor 
at  once. 

11.  Memorize    the    fact    that    acetic    acid    and    water    are    weak    electrolytes. 
Questions.     Refer  to  Assignment  21,  Paragraph  5,  and  state  what  you  would 
expect  to  happen  if  a  solution  containing  Ac~  at  high  concentration  were  mixed 
with  a  solution  containing  H+  at  high  concentration.     What  solutions  would  you 
mix  to  try  this  experiment,  and  how  could  you  prove  with  an  indicator  that  a 
reaction  had  taken  place? 

12.  The    reaction    between   sodium   acetate   and   hydrochloric    acid   solutions. 
Experiment.      Prepare    some    approximately    half    normal    hydrochloric    acid. 
Measure  out  two  15  cc.  portions  in  test-tubes,  and  add  two  or  three  drops  of 
methyl  violet  to  each.     Measure  10  cc.  1  N  sodium  acetate  and  add  the  solution, 

[24] 


a  few  drops  at  a  time,  to  one  of  the  0.5  N  hydrochloric  acid  solutions.  After 
each  addition  shake  the  mixture,  record  the  color  and  note  the  volume  of  the 
sodium  acetate  solution  added.  Write  the  equation  for  the  reaction  that  has 
taken  place  between  H+  and  Ac~.  Questions.  If  solutions  containing  0.10  mol 
HC1  and  0.05  mol  NaAc  were  mixed  what  substances  would  be  present  in  the 
resulting  solution?  What  would  be  the  concentration  of  each  if  the  final  volume 
were  (a)  1  liter,  (b)  500  cc.? 

13.  The  neutralization  of  NaOH  solution  by  acetic  acid.     The  concentration 
of  H+  in  an  acetic  acid  solution  is  small  in  comparison  with  the  total  concentration 
of  acid  in  the  solution,  but  the  student  must  not  jump  to  the  conclusion  that  the 
results  of  Assignments  5  and  6  on  neutralization   would  have  been  materially 
different  if  acetic  acid  had  been  used  throughout  instead  of  hydrochloric  acid. 
In  the  following  experiment  we  shall  study  qualitatively  the  reaction  between 
NaOH  solution  and  acetic  acid;  cf.  the  quantitative  experiment  in  Assignment  5, 
Paragraphs  4  and  5.     Experiment.     To  10  cc.  6  TV  NaOH  in  a  porcelain  dish 
add  about  8  cc.  6  Ar  HAc.     (Does  the  mixture  become  warm?)     Add  1  drop  of 
phenolphthalein  and  continue  to  add  acetic  acid  slowly  until  the  solution  becomes 
colorless;    test  again  with  the  indicator   (since  the  phenolphthalein  color  fades 
in  a  concentrated  NaOH  solution)  and  finally  add  about  2  cc.  acid  in  excess. 
Evaporate  the  solution  until  crystals  of  salt  begin  to  separate.     (Caution.     Do 
not  evaporate  the  solution  to  dryness,  since  the  NaAc  may  decompose.)     Allow 
the  mixture  to  cool,  collect  some  of  the  moist  salt,  NaAc  •  3  H2O,  and  dry  it 
between  filter  papers.     Prove  that  the  salt  contains  sodium  by  the  flame  test,  and 
acetate  by  warming  with  6  TV  sulfuric  acid  and  noting  the  odor  of  acetic  acid. 
Question.     What  evidence  is  furnished  by  this  experiment  that  the  neutralization 
reaction 

HAc  +  NaOH  ==    H2O  +  NaAc 

takes  place  when  solutions  of  HAc  and  NaOH  are  mixed?  The  concentration 
of  an  acetic  acid  solution  can  be  determined  by  titration  with  a  known  sodium 
hydroxide  solution  when  a  suitable  indicator  is  used,  phenolphthalein  in  this  case 
(and  if  time  permits  the  student  may  perform  this  titration). 

14.  We  shall  now  examine  the  above  reaction  in  the  same  way  as  we  have 
already  done  in  Assignment  21,  Paragraph  4,  in  the  case  of  the  reaction  between 
a  strong  acid  and  a  strong  base.     Questions.    Which  of  the  four  substances  are 
strong  and  which  weak  electrolytes?    What  substance  is  present  at  high  concen- 
tration in  an  acetic  acid  solution  which  is  not  present  at  high  concentration  in  a 
sodium  acetate  solution?     What  substance  is  present  in  a  NaOH  solution  which 
is  not  present  at  high  concentration  in  sodium  acetate  solution?  What  substance 
is  present  at  high  concentration  in  a  sodium  acetate  solution  which  is  not  present 
at  high  concentration  in  either  of  the  initial  solutions?     Write  the  equation  for 
the  main  reaction,  and  show  it  to  your  instructor.    The  study  of  this  reaction  will 
be  continued  in  Assignment  25.      V  <• 

15.  Problems,     (i)   20.0  cc.  acetic  acid  solution  were  found  to  neutralize  the 
same  volume  of  NaOH  solution  as  16.0  cc.  0.50  TV  H2SO4;  phenolphthalein  was 
the  indicator  in  both  titrations.     What  is  the  concentration  of  the  acetic  acid 
solution    (a)   in  mols  per  liter,   (b)   in  equivalents  per  liter,  and   (c)   in  grams 
per  liter? 

(2)   Outline  experiments  to  distinguish  between 

(a)  1.0  N  HNO3  and  0.10  N  HNO3 

(b)  A  solution  of  nitric  acid  and  one  of  acetic  acid  which  give  the  same 
bluish  color  with  methyl  violet. 


[25] 


ASSIGNMENT  23 

STRONG  AND  WEAK  BASES.     THE  USE  OF  INDICATORS  TO  MEASURE 
HYDROXIDE  ION  CONCENTRATION 


1.  In  Assignment  23  we  shall  develop  a  method  of  measuring  approximately,  by 
means  of  indicators,  concentrations  of  hydroxide  ion  between  normal  and  10~7 
normal,  the  concentration  in  pure  water.     This  method  will  be  used  in  studying 
solutions  of  a  typical  weak  base.     Since  there  is  throughout  a  close  relation  with 
the  preceding  Assignment  only  brief  direction  will  now  be  given.     The  student 
is  expected  to  make  use  of  the  discussion  and  Questions  in  Assignment  22  in 
correlating  the  results  of  the  two  Assignments. 

2.  Experiment.     Prepare  solutions  of  sodium  hydroxide  which  are  approxi- 
mately  normal,   0.1    normal,   0.01   normal,   and  0.001    normal.      (State   in  your 
note-book  how  you  prepared  these  solutions.)     To  10  cc.  of  each  solution  in  a 
test-tube,  add  1  drop  of  a  solution  of  the  indicator,  trinitrobenzol.     Record  the 
color  obtained  in  each  case,  and  observe  especially  the  most  dilute  solution  that 
gives  a  color  with  the  indicator.     Make  a  second  series  of  observations  using1  a 
larger  amount  of  the  indicator,  say  6  drops,  in  each  case.     Note  that  by  usirig 
the  different  amounts  of  indicator  you  can  determine  approximately  concentra- 
tions of  OH~  in  one  case  between  N  and  .01  N,  and  in  the  other  between  O.I  N 
and  0.001  N. 

3.  Repeat  the  experiment  with  potassium  hydroxide  solution,  and  compare  the 
colors    obtained    at    each    concentration.     If    sodium    hydroxide    in    solution    is 
completely   ionized,  what  conclusion  can  you   draw  with   respect  to  potassium 
hydroxide?     What  concentrations  of  hydroxide  ion  can  be  measured  by  means 
of  this  indicator,  trinitrobenzol? 

4.  Experiment.    Prepare  solutions  of  NaOH  or  KOH  which  are  approximately 
10-4  N  and   1O5  N.     Test  10  cc.  portions  of  the  10~3  N,   10~4  N  and   1Q-5  N 
solutions,  and  water  with  the  indicator  phenolphthalein.     Repeat,  using  litmus 
instead  of  phenolphthalein.    Note.    Since  the  quantity  of  alkali  in  a  given  volume 
of  these  dilute  solutions   is   extremely   small  it  is   evident  that  large   errors   in 
concentration  may  result  if  the  test-tubes  and  flasks  are  not  thoroughly  washed 
with  .distilled  water  before  use.     Check  your  results  by  preparing  fresh  portions 
of  the  10~3,  10~4  and  10~5  N  solutions  and  repeating  the  experiment. 

5.  Summarize  your  results  in  a  table  which  shows  the  color  obtained  with 
each  indicator  at  various  concentrations  of  hydroxide  ion.     Compare  this  table, 
and  the  corresponding  tph^e  in  Assignment  22,  Paragraph  6,  with  the  table  given 
by  Hildebrand  on  page  181. 

6.  In  the  next  experiment  we  shall  determine  the  concentration  of  hydroxide 
ion  in  solutions  of  the  weak  base,  ammonium  hydroxide,  NH^OH.     The  concen- 
trated laboratory  reagent,  cf.  Assignment  7,  Problem  (j),  is  prepared  by  dissolving 
the  gas  ammonia,  NH3,  in  water  until  the  solution  is  nearly  saturated  with  NH3 
at   room   temperature.      Write   the    equation    for   the    formation    of   ammonium 
hydroxide   from  ammonia  and  water.     Ammonium  sulfate,    (NH4).,SO4,  is  an 
example  of  a  salt  of  this  base.     Question.    What  is  the  concentration  in  mols  per 
liter  and  in  equivalents  per  liter  of  ammonium  ion,  NH4+,  and  of  sulfate  ion  in  a 
0.1  molal  solution  of  ammonium  sulfate? 

7.  Experiment.     From  the  laboratory  6  TV  NH4OH  solution  prepare  solutions 
which  are  approximately  N,  0.1  N,  0.01  N  and  0.001  N.    By  experiments  with  the 
indicators    trinitrobenzol    and    phenolphthalein    (write   out   the   details   of   these 
experiments  in  your  note  book)  estimate  the  concentration  of  OH~  in  at  least  two 
of  these  solutions.     In  each  case  give  also  the  concentration  of  NH4+  and  of 
un-ionized   NH4OH,  and  calculate  the  degree  of  ionization.     Note.     Accurate 

[261 


determinations  show  that  at  each  concentration  the  degree  of  ionization  of 
ammonium  hydroxide  is  almost  the  same  as  that  of  acetic  acid  at  the  same 
concentration;  see  Assignment  22,  Paragraph  10. 

8.  The  reaction  between  ammonium  chloride  and  sodium  hydroxide  solutions. 
Predict  what  will  happen  when  a  solution  of  ammonium  chloride  is  added  to  a 
solution  of  sodium  hydroxide;    cf.  Assignment  22,  Paragraphs  11  and  12.     Plan 
an  experiment  to  demonstrate  this  result  and  try  the  experiment.     Write  the 
ionic  equation  for  the  reaction. 

9.  The  neutralization  of  H^SO^  solution  by  ammonium  hydroxide  .  Experiment. 
To  10  cc.  6  N  H2SO4  in  a"  flask  add  6  N  NH4OH  from  a  graduate  until  the 
solution,  after  shaking,  has  a  distinct  odor  of  NH3.     Evaporate  in  a  porcelain 
dish  until  a  considerable  quantity  of  salt  has  separated,  cool  the  mixture,  collect 
some  of  the  salt  and  dry  it  between  filter  papers.     Prove  that  the  salt  contains 
(a)   sulfate,  by  dissolving  some  of  it  in  water  and  adding  barium  chloride   (to 
precipitate  BaSO4),  and   (b)   the  ammonium  radical,  by  warming  with  sodium 
hydroxide  solution  and  noting  the  odor  of  NH3.    Give  the  experimental  evidence 
and  reasoning  in  favor  of  the  conclusion  that,  when  solutions  of  a  strong  acid 
and  ammonium  hydroxide  solutions  are  mixed,  the  main  reaction  is: 

H+  +  NH4OH  =  H2O  +  NH4* 
The  study  of  this  reaction  will  be  continued  in  Assignment  25. 

10.  Problems.      (/)   State   how   you   would   determine   whether  an   unknown 
solution  is  more  acidic  or  more  basic  than  water.   (2)    Suggest  experiments  to 
distinguish  between   (a)   0.1  N  KOH  and  0.01  N  KOH.     (b)  A  solution  of  a 
strong   base    and    one    of   a    weak    base   which    have    the    same   hydroxide   ion 
concentration. 


ASSIGNMENT  24 
RAPID  REVERSIBLE  REACTIONS  AND  EQUILIBRIUM 


References.     Hildebrand,  Chapter  XII,  pages  155-171;  Chapter  XI,  pages  145 

and  148. 


1.  While   certain   reactions   proceed   to   completion,   as   the  transformation  of 
metallic  copper  into  cuprous  sulfide  studied  in  Assignment  1,  many  reactions  do 
not.    For  example,  when  solutions  containing  equivalent  amounts  of  a  strong  acid 
and  sodium  acetate  are  mixed,  cf.  Assignment  22,  Paragraph  12,  the  resulting 
solution  still  contains  about  1%  of  the  reacting  substances  H+  and  Ac~,  i.  e.,  the 
reaction  H+  -f-  Ac"  =  HAc,  although  it  takes  place  very  rapidly,  stops  when  about 
99%  of  the  possible  amount  of  HAc  has  been  formed.     In  the  final  solution,  the 
concentration  of  the  three  substances  involved  in  the  reaction  are  the  same  as  in 
the  corresponding  acetic  acid  solution.     Similarly,  if  we  had  started  with  pure 
acetic  acid,  which  is  un-ionized  in  the  solid,  liquid  or  gaseous  state,  and  dissolved 
it  in  water,  the  reaction  HAc  =  H+  +  Ac~  would  have  taken  place  rapidly,  but 
only  until  about  1%  of  the  acetic  acid  had  been  ionized.     Question.     From  your 
results  in  Assignment  22,  what  would  be  the  concentrations  of  H+,  Ac~  and  HAc 
in  a  solution  made  (a)  by  dissolving  1  mol  of  HAc  to  give  a  liter  of  solution,  and 
(b)  by  dissolving  1  mol  of  HC1  and  1  mol  of  NaAc  to  give  a  liter  of  solution? 

2.  The  reaction  just  considered  is  an  example  of  a  rapid,  reversible  reaction, 
and  the  three  substances  involved  in  this  reaction  are  in  equilibrium  with  each 
other  in  the  final  solution.     In  general  whenever  it  has  been  shown  experimentally 
that  a  reaction  can  be  made  to  take  place  in  both  directions,  i.  e.,  is  reversible, 
then  it  may  be  concluded  that  under  suitable  experimental  conditions  a  state  of 
equilibrium  can  be  realized  in  which  all  the  substances  involved  in  the  reaction 

[27] 


are  present  together.  For  each  set  of  experimental  conditions  there  is  a  definite 
state  of  equilibrium ;  and  when  the  experimental  conditions  are  altered,  e.  g.,  by 
changing  the  concentration  of  one  or  more  of  the  substances  involved  or  by 
changing  the  temperature  the  reaction  takes  place  in  one  direction  or  the  other 
until  equilibrium  is  again  established.  The  problem  is  to  learn  to  predict  what 
will  happen  in  any  given  case  when  the  experimental  conditions  are  altered. 

3.  The  effect  of  changing  the  concentration  of  one  of  the  substances  involved 
in  an  equilibrium.     Experiment.     Place  two  10  cc.  portions  of  N  acetic  acid  in 
two  test-tubes,  a  10  cc.  portion  of  0.1  N  acetic  acid  in  a  third  test-tube,  and  a 
10  cc.  portion  of  water  in  a  fourth  test-tube.     To  each  solution  add  the  same 
number  of  drops  of  methyl  violet  solution.    To  one  of  the  normal  solutions  add  a 
small  amount  of  solid  sodium  acetate  (or  of  4  AT"  solution)  ;  compare  the  colors 
of   the   four   solutions.     Repeat  the   experiment  with   more  dilute   solutions   of 
acetic   acid,   using  methyl   orange   instead   of   methyl   violet.      Question.     What 
conclusion   can  you   draw   with   regard  to  the   change   in  the   concentration   of 
hydrogen  ion  in  this  experiment?    What  reaction  must  have  taken  place?    When 
equilibrium  has  again  been  established  after  the  addition  of  the  sodium  acetate, 
is  the  concentration  of  each  of  the  substances  H+,  Ac"  and  HAc  greater  or  less 
than  its  concentration  in  the  original  acetic  acid  solution?     State  briefly  how  this 
experiment  illustrates  the  general  statement :    the  effect  of  changing  the  concen- 
tration of  one  of  the  substances  involved  in  an  equilibrium  is  to  cause  that  reaction 
to  take  place  which  tends  to  neutralize  the  change.     The  effect  of  increasing  the 
concentration  of  acetate  ion  could  also  have  been  predicted  from  the  quantitative 
statement  of  the  Mars  Law  :  (Concentration  of  H+)  (Concentration  of  Ac~)  /^Con- 
centration  of    HAc)  =  constant,    when   equilibrium   has   been   established   at   a 
definite  temperature. 

4.  Outline  an  experiment  to   demonstrate   that   the   reaction   NH4+-|-  OH~  = 
NH4OH  takes  place  when  a  solid  ammonium   salt  is  added  to   a   solution  of 
ammonium  hydroxide ;    cf .  your  experiments  in  Assignment  23.     Perform  this 
experiment,  and  explain  how  it  illustrates  the  italicized  statement  in  the  preceding 
Paragraph. 

5.  As  another  example  of  the  effect  of  change  of  concentration  upon  equilibrium 
we  shall  study  the  equilibrium  between  solid  silver  acetate  and  its  ions.     The 
reversible  reaction  is: 

AgAc  (  solid  )=  Ag+  +  Ac~ 

Note.  The  solubility  of  silver  acetate  is  0.06  mol  per  liter  at  room  temperature; 
while  it  is  much  more  soluble  than  a  sparingly  soluble  salt,  as  silver  chloride, 
it  is  much  less  soluble  than  such  salts  as  silver  nitrate,  sodium  nitrate  or  sodium 
acetate.  Experiment.  Prepare  some  solid  silver  acetate  by  adding  to  25  cc. 
4  N  NaAc  solution  (free  from  chloride)  *  about  20  cc.  0.1  N  AgNCX  solution, 
and  shaking  the  mixture  several  times.  Collect  the  solid  on  a  filter  paper,  and  dry 
it  by  pressing  between  dry  filter  papers.  Prepare  a  saturated  solution  by  shaking 
the  solid  with  10  cc.  water  at  intervals  for  about  10  minutes.  (The  saturated 
solution  can  be  prepared  somewhat  more  conveniently  by  warming  the  mixture  to 
40°  or  50°,  but  it  must  be  cooled  to  room  temperature  before  continuing  the 
experiment).  Allow  the  solid  to  settle.  Pour  half  of  the  clear  saturated 
solution  into  another  test-tube  and  set  it  aside  for  later  use,  Paragraph  7.  To  the 
remaining  mixture  of  solid  and  saturated  solution  add  6  N  HNO3  drop  by  drop, 
shaking  the  mixture  after  each  drop  is  added.  When  the  silver  acetate  has 
dissolved,  heat  the  solution  nearly  to  boiling  and  note  the  odor. 

Note.    Place  all  silver  residues,  including  any  solution  which  contain  an  appre- 
ciable amount  of  silver,  in  the  jar  maked  "silver  waste." 

*  If  chloride  is  present  it  may  be  removed  from  NaAc  solution  by  adding  first  some 
AgNOs  solution,  shaking  the  mixture  and  filtering  off  the  precipitated  AgCl.  The  resulting 
solution  contains  some  NaNOs  and  should  be  used  only  in  this  Assignment. 

[28] 


6.  The  disappearance  of  solid  silver  acetate  on  the  addition  of  a  strong  acid 
may  be  considered  to  be  due  to  a  shifting  of  the  equilibrium 

AgAc  (solid)  =  Ag+  +  A<r 

as  a  result  of  the  establishment  of  the  second  equilibrium 

H+  +  Ac-  =  HAc 

When  a  drop  of  nitric  acid  is  added  the  H+  uses  up  Ac~  to  form  HAc;  and 
in  order  to  again  build  up  the  concentration  of  Ac',  and  thus  restore  equilibrium, 
some  solid  AgAc  goes  into  solution.  When  more  HNO3  is  added  the  same 
processes  are  repeated,  until  finally  all  the  silver  acetate  is  dissolved.  From  this 
discussion  of  the  mechanism  of  the  reaction  it  is  evident  that  the  result  of  the 
experiment  could  have  been  predicted  from  the  fact  that  acetic  acid  is  a  weak 
acid.  The  principal  substances  that  have  disappeared  are  solid  AgAc  and  H+, 
and  those  that  have  formed  are  Ag+  and  HAc;  and  we  may,  therefore,  write 
for  the  main  reaction  the  equation : 

AgAc  (solid)  +  H+  ==  HAc  +  Ac-. 

The  same  equation  is  obtained  by  adding  together  the  two  equations  considered 
'  in  discussing  the  mechanism  of  the  reaction.  It  is  important,  however,  to  realize 
that  in  such  a  case  as  this  no  single  equation  can  represent  all  that  has  happened 
in  the  reaction.  Thus,  the  concentration  of  Ac~,  although  small  throughout  the 
experiment,  is  much  smaller  in  the  final  solution  than  in  the  initial  AgAc  solution ; 
this  small  decrease  is  due  to  the  reaction  H+  +  Ac~  =  HAc,  but  is  not  taken  into 
consideration  in  the  equation  which  we  have  written  for  the  main  reaction. 

7.  Predict  what  will  happen  when  solid  sodium  acetate  is  added  to  a  saturated 
solution  of  silver  acetate.     Experiment.     Test  your  answer  by  adding  a  small 
amount  of  4  AT  NaAc  to  the  clear  saturated  solution  of  silver  acetate  which  you 
have  just  prepared.     Question.    How  does  this  experiment  illustrate  the  italicized 
statement  in  Paragraph  3?     Re-word  that  statement  to  make  it  apply  specifically 
to  the  change  of  solubility  of  a  salt  when  the  concentration  of  one  of  the  ions  of  the 
salt  is  changed.    What  will  happen  when  solid  silver  nitrate  is  added  to  a  saturated 
solution  of  silver  acetate? 

8.  Problems.  (/)   For  each  of  the  following  reactions: 

(a)  H2O    (liquid)  =H2O  (gas) 

(b)  H2O   (solid)  =H2O  (liquid) 

(c)  NH3   (in  solution)  =  NH3  (gas) 

(d)  Ag+  +  Cl-  —  AgCl   (solid).   (Note  the  solubility  of  AgCl  is  1.0  X   10'5 
mols  per  liter)  ; 

outline  an  experiment  to  show  that  the  reaction  can  take  place  as  written  and  a 
second  experiment  to  show  that  the  reverse  reaction  can  also  take  place. 

(2)  In  each  case  in  Problem  (/)  point  out  the  conditions  under  which  equi- 
librium can  be  realized.  A  satisfactory  answer  in  (a)  is  the  following:  whenever 
water  and  water  vapor  are  present  together  in  a  closed  vessel  equilibrium  is 
quickly  established,  and  remains  unchanged  as  long  as  the  temperature  remains 
constant;  when  the  closed  vessel  contains  only  water  and  water  vapor  the 
pressure  at  equilibrium  is  1  atmosphere  when  the  temperature  is  100° ;  cf . 
Assignment  2,  Paragraph  10,  for  the  pressure  of  water  vapor  at  room  temperature. 

(j)  Give  an  example  of  a  reaction  which  can  be  shown  to  be  reversible  at  high 
temperature,  but  which  proceeds  so  slowly  at  room  temperature  that  equilibrium 
will  not  be  realized  in  several  months.  Note  that  all  the  substances  involved  in 
this  reaction  can  exist  together  at  room  temperature  but  that  such  a  fact  does  not 
constitute  a  proof  that  equilibrium  has  been  established. 

(4)   From  your  table  in  Assignment  22,  Paragraph  10,  state  what  fraction  of 

[29] 


the  acetic  acid  is  ionized  in  the  0.1  N  and  0.01  N  solutions.    What  reaction  takes 
place  when  1  liter  of  0.1  N  solution  is  diluted  to  10  liters  of  0.01  N  solution? 

(5)  State  how  the  experimental  result  in  Problem  (4)   could  have  been  pre- 
dicted  from   kinetic   considerations,   on  the   assumption  that  the  two   opposing 
reactions,  HAc  =  H+  -j-  Ac~  and  H+  -f-  Ac~  =  HAc,  are  actually  taking  place  at 
equilibrium. 

(6)  When  one  of  the  substances  involved  in  an  equilibrium  is  a  solid,  what 
is  the  effect  of  adding  more  of  this  solid  to  an  equilibrium  mixture?     Does  the 
solubility  of  a  salt  at  room  temperature  depend  on  the  amount  of  the  salt  in 
contact   with    the   saturated    solution?      Give    another   example    of    a   reversible 
reaction  involving  a  solid  substance. 


ASSIGNMENT  25 
THE  REVERSIBILITY  OF  NEUTRALIZATION  REACTIONS.    HYDROLYSIS 

References.     Hildebrand,  Chapter  XIII,  pages  185-192;  as  a  review,  read  pages 

173  and  180-183. 


1.  In  earlier  Assignments  we  have  studied  various  examples  of  neutralization 
reactions  and  have  shown  that  the  main  reactions  that  take  place  are  represented 
by  the  following  equations : 

(a)  H+  +  OH"  =  H2O  when  a  strong  acid  reacts  with  a  strong  base. 

(b)  HA  +  OH~  =  H2O  +  A~  when  a  weak  acid  reacts  with  a  strong  base. 

(c)  H+  +  BOH  =  H2O  +  B+  when  a  strong  acid  reacts  with  a  weak  base. 

In  the  present  Assignment  we  shall  find  that  these  reactions  are  reversible  and 
we  shall  study  the  equilibrium  that  is  established  in  each  case.  When  the 
Assignment  is  finished  the  student  should  be  able,  from  a  knowledge  of  the 
relative  strengths  of  various  acids,  and  of  various  bases,  to  predict  the  approxi- 
mate concentrations  of  the  various  substances  at  equilibrium  in  any  given 
neutralization  reaction. 

2.  We  wish  to  examine  experimentally  the  properties  of  the  solution  that  is 
obtained  when  equivalent  amounts  of  a  given  acid  and  base  react.    To  do  this  we 
shall  take  advantage  of  the  fact  that  a  solution  of  a  pure  salt  is  identical  with  the 
solution  prepared  by  mixing  exactly  equivalent  amounts  of  the  corresponding 
acid  and  base.     In  other  words  we  shall  make  use  of  the  reaction  which  is  the 
reverse  of  neutralization  in  establishing  the  equilibrium  state  that  we  desire  to 
study.    The  experimental  work  will  consist  in  determining  by  means  of  indicators 
the    approximate    concentration    of    H+    or   OH~    in    solutions   of   various    salts. 
Note.     In  the  following  experiments  it  is  important  that  all  test-tubes  should  be 
clean;    rinse  each  several  times  with  distilled  water  before  using  it. 

3.  Experiment.    Test  10  cc.  portions  of  N  NaCl  solution  and  distilled  water  by 
adding  to   each  the   same   number   of   drops   of  litmus   solution.     Repeat   using 
phenolphthalein  and  methyl  orange.     Question.     What  conclusions  can  you  draw 
with  regard  to  the  relative  concentration  of  H+  or  OH~  in  sodium  chloride  solution 
and  in  water?    Write  the  ionic  equation  for  the  neutralization  of  NaOH  by  HC1 
solutions,  and  state  whether  the  Na+  and  Cl"  are  involved  in  the  reaction.     In 
pure  water  concentrations  of  H+  and  OH~  are    the    same    and    equal    to    10~7 
mols  per  liter,   what   are  the  concentrations   of   Na+,   Cl~,   H+  and  OH~  in  the 
solution  made  (a)  by  dissolving  1  mol  NaCl  in  water  to  give  a  liter  of  solution, 
and  (b)  by  dissolving  exactly  1  mol  of  HC1  and  1  mol  of  NaOH  to  give  a  liter 
of  solution? 

4.  Experiment.     Test  10  cc.  portions  of  4  N  NaAc  and  distilled  water  first 

[301 


with  litmus  solution  and  then  with  phenolphthalein.  Estimate  approximately  the 
concentration  of  OH"  in  the  4  N  NaAc  solution.  Questions.  Is  the  concentration 
of  OH"  in  4  N  NaAc  solution  greater  or  less  than  in  pure  water?  From  the  fact 
that  the  solution  must  be  identical  with  that  obtained  by  mixing  exactly  equivalent 
amounts  of  acetic  acid  and  sodium  hydroxide,  what  conclusions  can  you  draw 
with  regard  to  the  completeness  of  the  neutralization  reaction  between  HAc  and 
OH"?  To  account  for  the  formation  of  OH"  when  NaAc  is  dissolved  in  water, 
what  reaction  must  take  place?  What  are  the  approximate  concentrations  of 
Na+,  Ac",  HAc  and  OH-  in  the  4  N  NaAc  solution? 

5.  It  will  be  observed  that  the  main  reaction  which  accounts  for  the  neutrali- 
zation of  HAc  by  NaOH, 

HAc  +  OH-  +  (Na+)  =  Ac-  +  H2O  +  (Na+) 

or  the  reverse  of  this  reaction  considered  in  the  preceding  paragraph,  involves 
the  disappearance  of  one  weak  electrolyte  and  the  formation  of  another.  Both 
the  OH"  and  the  Ac"  are  competing  for  the  H+.  From  the  fact  that  water  is  a 
far  weaker  electrolyte  than  acetic  acid,  it  could  have  been  predicted  that  when 
equilibrium  is  reached  the  OH"  will  have  taken  nearly  all  the  H+  away  from 
the  Ac" ;  i.  e.,  the  concentration  of  the  HAc  and  OH"  will  be  very  small  and  that 
of  the  Ac"  large.  Question.  From  your  table  of  the  concentrations  of  H+  in 
acetic  acid  solutions,  Assignment  22,  what  is  the  ratio  of  the  concentration  of  H+ 
in  N  acetic  acid  to  that  in  pure  water? 

6.  As  indicated  in  the  preceding  paragraph,  the  mechanism  of  this  reaction 
may  be  considered  to  be  parallel  to  that  of  the  dissolving  of  silver  acetate  in  a 
solution   of   a   strong  acid,   Assignment  24.     When   an   acetic   acid   solution  is 
neutralized  by  NaOH  solution,  the  OH"  unites  with  H+  which  is  present  at  small 
concentration  in  the  acetic  acid  solution : 

H+  +  OH-  =  H2O 

and,  in  order  to  again  restore  equilibrium,  the  HAc  dissociates, 

HAc  =  H+  +  Ac- 

These  two  reactions  continue  to  take  place  simultaneously  until  equilibrium  is 
established.  The  sum  of  the  two  gives  the  equation  which  we  wrote  for  the  main 
reaction. 

7.  Discuss  in  a  similar  way  the  mechanism  of  the  reaction  between  Ac"  and 
H2O  which  takes  place  when  sodium  acetate  is  dissolved  in  water.     This,  of 
course,  involves  the  reverse  of  the  two  reactions  written  in  Paragraph  6.     Write 
the  equation  for  the  main  reaction  in  this  case. 

8.  The   reaction  which   is   the   reverse  of  neutralization  is   called  hydrolysis. 
In  the  reaction  which  we  have  just  considered  a  small  amount  of  the  acetate  ion 
is   said  to   be   hydrolyzed.     Question.     Why  would  you   expect   the  Ac"  to  be 
hydrolyzed  slightly  while  the  Cl~  is  not  hydrolyzed  at  all? 

9.  The   dissociation  of   HCXV,   bicarbonate   ion,   into  H+   and   carbonate  ion, 
CO3",  is  far  less  than  the  dissociation  of  acetic  acid.     Questions.    What  reaction 
would  you   expect   to   take  place   between   CO3~~  and   H2O   when   Na2CO3   is 
dissolved  in  water?    Which  would  you  expect  to  be  the  more  basic,  a  solution  of 
Na2CO3  or  a  solution  of  NaAc  of  the  same  molal  concentration?     Test  your 
answer  by  the  following  experiment 

10.  Experiment.     Estimate  the  approximate  concentration  of  OH"  in  a  0.5  M 
Na2CO3    solution   by   testing   10   cc.   portions   with   litmus,   phenolphthalein   and 
trinitrobenzol.     The  main  reaction  for  the  hydrolysis  of  CO3~  "  is 

C03--  +  H20  =  HCO8-  +  OH- 

and  the  concentration  of  OH~  in  a  0.5  M  Na2CO3  solution  is  about  0.01  M. 
Questions.  What  would  be  the  concentration  of  OH-  if  the  CO3-  "  were  completely 

[31] 


hydrolyzed?  What  fraction  of  the  CO3~~  has  been  hydrolyzed?  What  would  be 
the  approximate  concentrations  in  mols  per  liter  of  Na+,  CO3~~,  HCO3~  and  OH" 
in  a  solution  prepared  by  dissolving  0.05  mol  NaHCO3  and  0.5  mol  NaOH  to 
give  1  liter  of  solution? 

11.  Experiment.     Test  a  10  cc.  portion  of  4  AT  NH4C1  with  litmus  solution. 
Compare  the   color  with  that  obtained  by  adding  the  same  amount  of  litmus 
solution  to  (a)  10  cc.  water  and  (b)  10  cc.  water  containing  1  drop  of  6  N  HC1. 
Questions.    How  do  you  account  for  the  fact  that  a  solution  of  NH4C1  is  slightly 
acid  while  that   of  NaAc  is   slightly  basic.     Write  the   main   reaction   for  the 
hydrolysis  of  the  NH4+,  and  the  two  reactions  which  may  be  used  to  explain  the 
mechanism  of  the  reaction.     The  concentration  of  H+  in  4  N  NH4C1  is  about 
4  X  10~5   N.     What   is   the   concentration   of   each   of  the  principal   substances 
present  in  a  solution  made  by  dissolving  4  mols  of  HC1  and  4  mols  of  NH4OH 
to  form  a  liter  of  solution? 

12.  Problems.      (1)   What     is     the     concentration     of     H+     in     N     NaOH? 
Note.     Consult  Hildebrand,  page  180. 

(2)  From  the  results  of  the  experiments  in  this  Assignment  and  the  table  of 
indicators  (Hildebrand,  page  181)  select  one  or  more  indicators  which  might  be 
used  to  determine  the  end-point  (i  e.,  to  determine  when  equivalent  amounts  of 
acid  and  base  are  present)  in  each  of  the  following  titrations : 

(a)  NaOH  solution  with  HC1  or  HNO3  solution. 

(b)  HAc  solution  with  NaOH  solution. 

(c)  NaHCO3  solution  with  NaOH  solution. 

(d)  NH4OH  solution  with  HC1  solution. 
Note.     Compare  Assignment  6,  Paragraph  3. 

(5)  An  ammonium  acetate  solution  gives  the  same  colors  with  indicators  as 
does  water :  What  hydrolysis  reactions  must  occur  when  solid  NH4Ac  is  dissolved 
in  water?  State  what  substances  are  present  in  the  solution. 

(4)  Point  out  parallelisms  between  the  strengths  of  acids  and  bases  and  the 
hydrolysis  of  the  corresponding  ions. 


32 


SECTION  III 
REACTIONS  OF  IONS 


ASSIGNMENT  31 

THE  PROPERTIES  OF  SODIUM,  POTASSIUM  AND  AMMONIUM  IONS. 
TESTS  FOR  CHLORIDE,  SULFATE  AND  NITRATE  IONS 


Reference,  for  Assignments  31  and  all  later  Assignments:    A  Standard  Text  on 

Inorganic  Chemistry. 

1.  In  this  Assignment  we  shall  study  the  properties  of  the  ions  of  the  common 
acids  and  bases. 

2.  Aside  from  the  physical  properties,  such  as  color  and  taste,  the  properties 
of  an  ion  depend  on  its  behavior  toward  other  substances.     The  study  of  the 
chemistry  of  an  ion  thus  consists  in  determining  whether  or  not  various  substances 
react  with  the  ion,  and  in  studying  the  reactions  that  do  take  place.     For  each 
reaction  it  is  necessary  to  know : 

(1)  the  products  of  the  reaction, 

(2)  whether  the  reaction  takes  place  rapidly  or  slowly, 

(j)   whether  the  reverse  reaction  can  be  made  to  take  place,  and 

(4)   the  nature  of  the  equilibrium  if  the  reaction  is  rapid  and  reversible. 

In  this  course  the  task  is  greatly  simplified  on  account  of  the  following  consid- 
erations :  Nearly  all  the  reactions  studied  take  place  very  rapidly ;  this  is 
characteristic  of  ionic  reactions,  such  as  ionization,  precipitation,  neutralization, 
etc.  Also  the  number  of  different  types  of  reactions  is  not  large ;  and,  when  for 
one  type  a  single  example  is  understood,  additional  examples  should  present  no 
real  difficulty.  Finally  an  equilibrium  state  is  reached  in  many  cases,  and  a 
knowledge  of  the  principles  of  equilibrium  enables  the  student  to  predict  what 
will  happen  when  there  is  a  given  change  in  the  experimental  conditions.  While 
the  number  of  facts  to  be  remembered  is  thus  greatly  reduced,  it  is  necssary, 
however,  to  memorize  such  facts  as;  the  formulas  of  substances,  the  relative 
solubility  of  salts,  the  relative  volatility  of  various  solid  or  liquid  substances  and 
the  relative  strength  of  weak  electrolytes. 

3.  Properties  of  sodium  ion.     The  chemistry  of  Na+  is  extremely  simple  and 
may    be    summarized    by    saying    that,    at    ordinary    temperatures,    with    a    few 
uncommon  exceptions,  all  of  its  compounds  are  non-volatile,  readily  soluble  and 
strong  electrolytes. 

4.  Flame  test  for  sodium.     Experiment.     Clean  an  iron  wire  by  dipping  it  into 
dilute  HC1  solution  in  a  test-tube  and  holding  the  wire  in  the  flame  until  it  gives 
only  a  faint  yellow  color.      (Note.     Do  not  contaminate  your  HC1  solution  hy 
dipping  the  iron  wire  into  the  bottle.)     Try  the  flame  test  with  a  solid  sodium  salt, 
the   laboratory    solution   of   a   sodium   salt,   a  very   dilute   solution  and   distilled 
water.      The   test   is   extremely   delicate   and   substances   which   do   not   contain 
measurable  amounts  of  sodium  usually  give  a  slight  yellow  color  for   a  short 
time.     Therefore,  before   concluding  that  sodium  is  present  in  an  "unknown" 
it  is  well  to  make  comparative  tests  with  known  solutions  (/)  free  from  sodium 
and    (2)    the  same  solution   containing  a  small   amount  of   sodium  salt.     Such 
experiments  are  called  blank  tests. 

5.  Properties  of  potassium  ion.     What  is  the   relation  of  K  and  Na  in  the 
Periodic    System?      Potassium   ion    resembles  .sodium   ion   very   closely.       One 
difference,  however,  is  the  precipitation  of  potassium  cobaltinitrite  when  cobalti- 
nitrite  ion,  Co(NO2)fi is  added  to  a  solution  containing  K+.    Experiments.    To 

[  33  ] 


5  cc.  of  a  dilute,  neutral  solution  of  a  potassium  salt  add  a  few  drops  of  acetic 
acid  and  2  to  5  cc.  of  the  sodium  cobaltinitrite  reagent ;    let  the  mixture  stand 
10  minutes. 

6.  Flame  test  for  potassium.    Try  the  flame  test  with  solid  KC1,  KC1  solution, 
KC1  solution  containing  a  small  amount  of  Nad,  Nad  solution,  and  NaCl  solution 
containing  a  small  amount  of  KC1.     Use  the  blue  glasses  to  shut  off  the  yellow 
color  of  the  sodium  flame  and  practice  until  the  presence  or  absence  of  both 
potassium  and  sodium  can  be  determined.     Note.    Since  the  flame  test  depends 
upon  the  amount  of  solid  present  it  is  well  to  concentrate  a  dilute  solution  by 
evaporation  before  making  the  test.     Chlorides  give  more  satisfactory  flame  tests 
than  sulfates  or  oxides,  which  is  due  to  the  relatively  greater  volatility  of  the 
chlorides. 

7.  Properties  of  ammonium  Ion.    The  ammonium  ion  resembles  sodium  ion  and 
potassium   ion  in   that  with   few  exceptions   its   compounds  are   soluble,  strong 
electrolytes.     Like  the  potassium  ion  it  forms  a  precipitate  with  Co(NO2)6       . 
Repeat  the  experiment  in  Paragraph  4,  using  NH4+  instead  of  K+.     Ammonium 
compounds  give  no  flame  test.     The  chemistry  of  NH4+  is  made  somewhat  more 
complicated  by  the  fact  that  NH4OH  not  only  is  a  weak  base,  but  is  decomposed 
easily  into   NH3  and  H2O.     NH3,  ammonia,  enters  into  many  reactions  which 
will  be  considered  in  later  Assignments.     Experiment.     Note  the  odor  of  the 

6  N  NH4OH  laboratory  reagent;  prepare  some  dilute  solutions  of  NH4OH,  e.  g., 
0.1  A7"  etc.,  heat  gently  about  10  cc.  of  each  solution  (moving  the  test-tube  back 
and  forth  through  the  flame)  and  note  the  odor  from  time  to  time;   decide  what 
is  the  most  dilute  solution  you  can  detect  by  means  of  the  odor  of  NH3  gas. 
Determine  by  experiment  whether  you  can  detect  the  odor  of  NH3  gas  when  a 
solution  of  NH4C1  or   (NH4)2SO4  is  boiled.     Suggest  a  method  of  testing  for 
ammonium  ion  based  on  the  formation  first  of  NH4OH  in  the  solution  and  then 
of  NH3  gas.     Write  equations.     Experiment.     Try  your  method  with  a  dilute 
solution  of  an  ammonium  salt. 

8.  Volatility    of   ammonium   compounds.      Ammonium   compounds   are   much 
more  volatile  than  those  of  sodium  and  potassium.     Experiment.     Evaporate  a 
small  amount  of   (a)   NH4C1  solution  and   (b)   a  mixture  of  NH4C1  and  KC1 
solutions  to  dryness  in  a  porcelain  dish  and  heat  until  fumes  are  no  longer  given 
off.     Suggest  a  method  other  than  the  flame  test  of  determining  the  presence  of 
K+  in  a  solution  which  contains  both  K+  and  -NH4+. 

9.  Properties  of  H+and  OH~.     Aside    from    the  tendency    of  H+  and  OH"  to 
unite  with  each  other  to  form  water  each  of  these  ions  reacts  with  many  other 
ions  to  form  weak  electrolytes.    We  have  already  noted  the  reaction  between  H+ 
and  Ac~  and  OH~  and  NH4+.  Acids  are  more  volatile  than  their  salts :  the  relative 
volatility  of  the  common  acids  will  be  considered  in  Assignment  34.     The  OH~ 
also   forms  precipitates  with  many  metallic  ions.      Question.     How  would  you 
test  for  (and  estimate  the  concentration  of)  H+  or  OH"  in  an  unknown  solution? 

10.  Although  the  ions  NO8~,  Cl~  and  SO4"  will  not  be  considered  in  detail  in 
this  Assignment,   certain  of  their  reactions  will  be  studied  in  order  that  their 
presence   in   solutions   may   be   detected.      All   nitrates   are   soluble,   but   a   few 
chlorides  and  sulfates  are  sparingly  soluble. 

11.  Test  for  chloride  ion.    Experiment.    To  1  cc.  portions  of  solutions  of  NaCl, 
Ba(OH)2,  Na2CO3  and  Na2SO4  add  10  cc.  water,  shake  and  add  a  few  drops 
of  AgNO3  solution.     In  each  case,  record  your  observations,  noting  the  color  of 
the  precipitates.     Now  add  HNO3  until  each  solution  is  acid.    From  your  results 
state  how  you  would  test  an  unknown  solution  for  chloride  ion.    Note.   Ba(OH)2 
was  used  instead  of  NaOH  because  Cl~  is  usually  present  in  the  laboratory  NaOH 
solutions.    Experiment.     Test  a  dilute  solution  of  your  NaOH  for  Cl~. 

[34] 


12.  Test  for  Sulfate  ion.     Experiment.     To  dilute  solutions  of  NaCl,  NaOH, 
NaXO3  and  Na2SO4  add  a  few  drops  of  Ba(NO3)2  solution.  Now  add  HNO3 
until  each  solution  is  distinctly  acid.     From  your  results  state  how  you  would 
test  an  unknown  solution  for  SO4~~. 

13.  Test  for  nitrate  ion.     Try  the  following  experiment  with  a  solution  which 
contains  nitrate  ion  and  with   one   free   from  nitrate  ion,   first   acidifying  with 
6  N  H2SO4  if  the  solution  is  alkaline.     To  about  2  cc.  of  the  solution  to  be 
tested  for  nitrate  ion  add  in  excess  a  solution  of  ferrous  sulfate,  FeSO4  (or  of 
ferrous   ammonium   sulfate, — a  double  salt  of   ferrous   sulfate   and   ammonium 
sulfate),  filter  if  there  is  a  precipitate,  hold  the  test  tube  in  a  slanting  position, 
and  pour  carefully  down  the  side  of  the  test-tube  (from  a  small  beaker)   2  or 
3   cc.   concentrated   sulfuric  acid.     The   concentrated  acid   sinks   to   the  bottom 
of  the  tube  and  a  dark  brown  ring  forms  on  its  surface  when  nitrate  is  present. 
Note.      The    brown    substance    decomposes    rapidly    when    the    mixture    is   hot. 
If  the  result  of  the  test  is  negative  and  the  test-tube  feels  hot  to  the  hand  (owing 
to  the  heat  liberated  when  the  concentrated  H2SO4  mixes  with  the  solution), 
repeat  the  test  more  carefully. 

14.  Make  up  solutions  of  various  concentrations  of  Cl~  and  try  experiments 
to  determine  if  your  test  can  be  used  to  decide  whether  the  concentration  of 
Cl~  is  large  or  small.     Repeat  for  the  nitrate  and  sulfate  tests. 

15.  Analyses  Nos.  A  and  B.     Analyze  the  unknowns  for  the  four  positive  and 
four  negative   constituents  considered  in   this  Assignment.     Try  to  distinguish 
between  large  amounts,  small  amounts,  and  traces. 

16.  Problems.    (/)  What  is  the  valence  of  each  of  the  4  positive  and  4  negative 
ions  considered  in  this  Assignment?    Write  the  formulas  of  the  sixteen  possible 
compounds,   each  of  which  contains  one  positive  and  one  negative  constituent 
(cf.    Hildebrand,   pages   90-91).      Name   each   of   the   sixteen   substances,   state 
whether  it  is  a  solid,  liquid  or  gas  at  room  temperature,  whether  it  is  readily 
soluble  in  water  or  not,  and  whether  it  is  a  strong  or  weak  electrolyte. 

(2)  What  is  the  valence  of  sodium  in  (a)  solid  sodium  sulfate,  (b)  metallic 
sodium? 

(j)  A  solid  unknown  was  dissolved  in  wafer  and  the  solution  was  found  to  give 
distinct  tests  for  Na+,  K+,  Cl~  and  SO4~  ~.  What  conclusions  can  you  draw  with 
regard  to  the  nature  of  the  solid  salts  in  the  original  mixture? 

(4)  Outline   experiments    to   decide   whether   or   not   each   of   the   following 
contains  the  impurity  named  : 

(a)  Barium  chloride  impurity  in  the  barium  hydroxide  solution. 

(b)  Potassium  chloride  in  solid  sodium  chloride. 

(c)  Potassium  chloride  in  solid  ammonium  chloride. 

(d)  Ammonium  chloride  in  solid  potassium  chloride. 

(e)  Sodium  nitrate  in  solid  sodium  sulfate. 

(5)  In  each  of  the  following  it  is  assumed  that  two  solutions  are  mixed  each 
of  which  contains  one  of  the  substances  at  moderate  concentration,  say  0.1  M . 
Mark  the  cases  in  which  no  reaction  takes  place.    In  the  others  write  the  equation 
for  the  reaction.     In   (a),   (b),   (c)   state  what  solutions  you  would  use  in  the 
experiment. 

(a)  Ac-andCl-  (d)  NH4OH  and  K*  (g)   K+  and  Ac~ 

(b)  NHt+  and  Na+  (e)   H+ and  Ac-  (h)   Ag+  and  Cl- 

(c)  NH4OHandH+  (/)   K+ and  OH-  (i)   Ba++ and  Cl~ 

F35  1 


ASSIGNMENT  32 
CALCIUM  ION 


References.  Hildebrand,  pages  175-179,  192-193  and  207. 


1.  In  Assignment  32  we  shall  study  the  chemistry  of  metallic  calcium  and  of 
calcium  ion,  including  its  formation  from  metallic  calcium.     We  shall  find  that 
certain  compounds  of  this  element  differ  from  the  corresponding  compounds  of 
sodium,  potassium,  and  ammonium  in  that  they  dissolve  to  a  much  smaller  extent 
in  water;  and  we  shall  study  the   equilibrium  between  these  solids  and  their 
saturated  solutions. 

2.  Note.    In  future  write  equations  for  all  reactions. 

3.  Obtain  from  the,  office  a  piece  of  metallic  calcium.    Describe  its  properties  as 
far  as  you  can  observe  them  by  physical  examination.     In  what  respects  does 
calcium  show  the  physical  properties  of  a  metal? 

4.  Reaction    between    calcium    and   water.      Experiment.       Drop  the  calcium 
into  about  20  cc.  of  water.     Stir  the  mixture  and  warm  it  gently  until  the  metal 
has  disappeared.    Test  the  solution  for  OH-.    What  reaction  has  occurred  ?   The 
white  solid  formed,   calcium  hydroxide,   Ca(OH)2,   is   a  strong  base  but  only 
moderately  soluble  in  water.     Cork  the  test-tube  and  save  the  mixture  for  use 
in  a  later  experiment.     Question.  Name  two  other  metals  that  react  readily  with 
water.     What   reactions   occur? 

5.  Calcium  hydroxide.     If  solid  calcium  hydroxide  is  shaken  with  water  until 
a  saturated  solution  is  obtained  an  equilibrium  between  solid  Ca(OH)2  and  the 
ions  Ca+*  and  OH~  is  established.    The  equation  for  the  reaction  that  has  taken 
place  is 

Ca  (OH  )2  (solid)  =  Ca++  +  2OH- 

The  solubility  of  Ca(OH)2  at  room  temperature  is  0.02  mols  per  liter.  Questions 
What  is  the  concentration  of  calcium  ion,  of  hydroxide  ion,  (/)  in  mols  per 
liter,  and  (2)  in  equivalents  per  liter?  Is  the  reaction  just  considered  reversible? 
What  solutions  would  you  mix  in  order  to  prepare  solid  Ca(OH)2?  Try  the 
experiment. 

6.  Which  contains  the  larger  concentration  of  OH~,  a  saturated  solution  of 
Ca(OH)2  or  1  AT  NH4OH   (see  Assignment  23)?     State  what  you  think  will 
happen  when  a  dilute  solution  of  calcium  chloride  is  made  alkaline  with  ammonium 
hydroxide.     Try   the   experiment.     Explain   the  result   if   your  prediction   was 
incorrect. 

7.  State  what  would  happen  if  a  solution  of  calcium  hydroxide  were  treated 
with  (i)    hydrochloric    acid,   (2)   ammonium    chloride.       Predict    what    would 
happen  if  solid  Ca(OH)2  were  also  present  in  each   case.  Experiment.   Test 
your  answers   by  experiments  with  portions  of   the  mixture  prepared   in   Par- 
agraph 4.  Questions.  How  would  you  use  calcium  hydroxide  to  test  a  solution  for 
ammonium  ion? 

8.  Calcium  carbonate.    The  solubility  of  CaCO3  in  water  Lt  room  temperature 
is  very  small,  viz.,  0.00013  mols  per  liter.     Question.     What  is  the  concentration 
of  calcium  ion,  of  carbonate  ion,   in   this  saturated  solution,    (i)    in  mols  per 
liter,   (2)  in  equivalents  per  liter?     What  solutions  would  you  mix  in  order  to 
form    a   precipitate    of    calcium    carbonate?    Try    the    experiment.      Write    the 
simplest  ionic  equation   for  the  reaction.     Continue  the  experiment  with  very 
dilute  solutions  of  a  calcium  salt  in  order  to  determine  if  this  reaction  can  be 
used  as  a  delicate  test  for  calcium  ion.     In  the  more  dilute  solutions,  if  a  pre- 
cipitate does  not  form  at  once,  heat  to  boiling  and  let  stand  10  minutes.     If  the 
liquid  is  turbid  compared  with  water  a  precipitate  has  formed. 

[361 


9.  Experiment.     Dilute  5  cc.  N  CaCL  with  about  100  cc.  water,  add  about  8cc. 
N  Na2CO3  solution,  heat  the  mixture  to  boiling,  filter,  wash  the  precipitate  and 
reject  the  wash  water.     Consider  what  substances  may  be  present  in  the  filtrate. 
Test  for  chloride  ion  and  for  calcium  ion.    Which  reagent  was  present  in  excess, 
calcium  chloride,  or  sodium  carbonate? 

10.  Experiment.     Treat  a  portion  of  the  CaCO3  precipitate  with  hydrochloric 
acid    solution.        Repeat  the    experiment    with    nitric    acid.        The    gas    given 
off  is  carbon  dioxide,  CO2.    What  substances  are  present  in  the  final  solution  in 
each  case  after  the  solution  has  been  boiled  to  expel  the  CO2?    Note.     CO2  is 
moderately    soluble    in    cold   water    and    the    solution    contains    carbonic    acid, 
H2CO3,  which  is  a  weak  acid.     CO2  gas  will  not  be  evolved  until  the  concentra- 
tion of  carbonic  acid  in  the  solution  becomes  sufficiently  high.     Questions.     How 
is  the  equilibrium  between  solid  CaCO3  and  its  ions  disturbed  by  the  addition 
of  a  strong  acid?     What  is  the  analogy  between  this  example  and  the  action 
of  an  acid  on  solid  Ca(OH)2,  and  on  solid  silver  acetate?     What  is  the  reason 
for  the   difference  between  these   results .  and  that  obtained  when  dilute  nitric 
acid  is  added  to  silver  chloride  or  barium  sulfate? 

11.  Test  for  carbonate.    Experiment.    Fit  a  test-tube  with  stopper  and  delivery 
tube.     Place  a  small  amount  of  any  carbonate  in  the  tube  and  add  a  few  cc.  of 
6  N  H2SO4.    Insert  the  stopper  and  pass  the  gas  given  off  into  a  clear  solution  of 
Ca(OH)2.*  If  only  a  small  amount  of  carbonate  is  present  the  CXX  will  not  pass 
over  into  the  Ca(OH)2  solution.  A  satisfactory  test  may  be  obtained  in  this  case 
by  placing  the  test-tube  in  a  beaker  of  boiling  water,  or  by  adding  a  small  piece  of 
zinc  to  the  cold  mixture  in  the  test-tube.     The  CO2  is  carried  over  with  the 
dissolved  air  and  steam  in  the  first  method,  and  with  the  hydrogen  gas  in  the 
second. 

12.  Calcium    oxalate.      Experiment.      Prepare    some    CaC2O4    by   treating   a 
solution   containing   calcium   ion   with   ammonium   oxalate   solution.      Test   the 
action  of  excess  of  a  strong  acid,  as  HC1  or  HNO3  on  CaC2O4.     Question.     Is 
oxalic  acid,  H2C2O4,  a  weak  or  a  strong  acid? 

13.  Experiment.     Determine    if    the    precipitation    of    CaC2O4    is    a    delicate 
test  for  calcium  by  treating  very  dilute  solutions  of  Ca++  with  ammonium  oxalate 
solution  and  a  little  ammonium  hydroxide  (to  make  sure  that  the  mixture  is  not 
acid).     Observe  the  same  precautions  as  in  the  test  for  calcium  by  means  of 
CO3    .    Note.    The  solubility  of  calcium  oxalate  in  water  is  even  less  than  that 
of   calcium   carbonate. 

14.  Calcium   sulfate.     Experiment.     To    10   cc.    normal   CaCl2   solution   add 
2  cc.   6  normal  H2SO4.     If  no  precipitate  appears  at  once,  heat  the  solution 
gently,  and  let  it  stand.    Filter.    Test  a  portion  of  the  filtrate  for  Ca++  by  adding 
NH4OH  until  the  solution  is  no  longer  acid,  and  then  (NH4)2CO3  solution,  and 
warming  the  mixture.    Test  another  portion  for  SO4~~.    What  conclusion  do  you 
draw  with   regard   to   the   solubility   of   CaSO4  in   water?     Which   is   the  less 
soluble,  (i)  CaSO4  or  CaCO3,  (2)  CaSO4  or  BaSO4? 

15.  Predict  what   will  take   place   when   solid   CaSO4   is   heated   with   excess 
Na2CO3   solution?     Experiment.     Test  your  prediction  by  boiling  some  solid 
CaSO4  with  normal  Na2COs  solution.     Test  the  filtrate  for  SO4~.     Wash  the 
precipitate  with  water,  and  test  it  for  carbonate. 

16.  Experiment.    Test  whether  the  reverse  reaction  will  take  place,  by  heating 
solid  CaCO3  with  sodium  sulfate  solution,  filtering  and  testing  the  precipitate  and 
filtrate. 

^Instead  of   Ca(OH)2  solution,   Ba(OH)s.  solution  may  be  used;    BaCO3,   like   CaCOs  is 
a  difficultly  soluble  substance. 

[37] 


17.  Flame  test  for  calcium.     Experiment.     Determine  whether  calcium  salts 
give  a  characteristic  flame  test.     Can  you  distinguish  the  flame  with  a  calcium 
salt  from  that  with  sodium  or  potassium  salts? 

18.  Note.     The  solubilities  of  many  common  salts  at  18°  are  listed  on  the  inside 
cover    page    of    Alexander    Smith's    text-book.      Do    not    try   to    remember   the 
actual  solubilities,  but  make  lists  of  the  readily  soluble  and  difficultly  soluble 
salts  of  each  metal  studied  and  memorize  these  lists.  Arrange  the  compounds  ot 
calcium  according  to  their  solubilities  in  water,  distinguishing  readily  soluble, 
moderately  soluble,  and  difficultly  soluble  substances.     Point  out  the  compounds 
which  are  much  more  soluble  in  dilute  hydrochloric  or  nitric  acids  than  in  water. 

19.  Give    as    many    methods   as   you   can    of   testing    for   Ca++   in   a   solution 
which  is  known  to  contain  H+,  Na+,  K+  and  NH4+. 

20.  Problems,    (i)     What  is  the  minimum  volume  of  water  needed  at  18°  to 
dissolve  i  gram  of  calcium  carbonate? 

(2)  Can  the  following  substances  be  present  at  moderate  concentrations  in 
the  same  solution?  If  not,  what  is  formed? 

(a)   H+andNO3-  (e)   H+  and  C2O4-  - 

(0)   H+  and  OH-  (/)   Ca++  and  CO3-  - 

(c)  H+  and  SO4-  -  (g)   Ca++  and  NO3~ 

(rf)   H+  and  CO3-  -  (h)  Ca++  and  NH4OH 

(j)   How  would  you  prepare: 

(a)  Solid  CaCO3  from  solid  Ca(OH)0? 

(b)  Solid  CaSO4  from  solid  CaCO3?~ 

(c)  Solid  CaC2Ot  from  solid  CaSO4? 

Write  equations  for  all  reactions,  and  point  out  what  equilibria  are  involved. 

(4)  What  is  the  position  of  calcium  in  the  Periodic  System?  Name  the 
elements  in  the  alkaline  earth  group.  Write  a  brief  note  illustrating  gradations 
of  properties  in  this  group,  after  examining  the  table  given  by  Hildebrand  on 
page  268. 


ASSIGNMENT  33 
CARBONATE  ION,  BICARBONATE  ION,  AND  CARBONIC  ACID 


1.  This  assignment,  in  which  we  shall  study  the  chemistry  of  carbonic  acid  and 
its  ions,  is  introduced,  not  only  because  this  subject  is  an  extremely  important 
one,  but  also  because  it  serves  to  illustrate  how  a  large  number  of  facts  can  be 
correlated  when  the  equilibria  involved  in  a  few  simple,  reversible  reactions  are 
understood.    Carbonic  acid,  H2CO3,  is  a  weak,  dibasic  acid;  and,  like  other  weak 
acids  which  contain  more  than  one  acidic  hydrogen  in  the  molecule,  it  ionizes  in 
steps,  to  form  (/)  bicarbonate  ion,  HCO3~,  and  (2)  carbonate  ion,  CO3~~,  thus: 

H.CO,  ==  H+  +  HCCV  (1) 

HCO3-  =  H+  +  CO3--  (2) 

In  order  to  understand  what  happens  in  the  various  reactions  considered,  it  is 
necessary  to  know  the  relative  concentrations,  in  the  initial  and  final  solutions, 
of  the  substances  involved  in  these  two  rapid  reversible  reactions.  Our  problem, 
then,  is  to  determine  how  the  concentrations  of  H+  (or  OH"),  H2COo,  HCO3~ 
and  CO3~~  at  equilibrium  in  a  solution  vary  as  the  experimental  conditions  are 
changed. 

2.  Obtain  at  the  office  a  large  and  small  two-holed  rubber  stopper,  two  pieces 
of  rubber  tubing  and  a  thistle  tube  for  use  in  the  experiment  in  paragraph  3  arid 
a  hard  glass  test-tube  for  use  in  the  experiment  in  Paragraph  12. 

[381 


3.  Preparation  and  properties  of  carbon  dioxide.    Experiment.    Make  a  carbon 
dioxide     generator     by     fitting    your     half-liter     flask     with     the     thistle     tube 
(passing  through  the  stopper  nearly  to  the  bottom  of  the  flask),  and  an  outlet 
tube  bent  at  right  angles.     Make  a  wash  bottle  for  washing  the  CO2  gas  by 
equipping  a  small  flask  with  an  entry  tube  extending  nearly  to  the  bottom  of  the 
flask,  and  a  delivery  tube.     Fill  the  small  flask  about  half  full  of  water.     Place 
in  the  generator  a  few  lumps  (about  5  grams)  of  limestone,  cover  with  water, 
and  add  hydrochloric  acid,  a  little  at  a  time,  through  the  thistle  tube. 

Note  the  color  and  the  odor  of  the  gas.  Calculate  from  the  molecular  weights 
the  relative  densities  of  carbon  dioxide  and  oxygen,  and  of  carbon  dioxide  and, 
nitrogen,  at  the  same  temperature  and  pressure.  Question.  Is  carbon  dioxide 
denser  or  lighter  than  air?  Experiment.  Collect  some  carbon  dioxide  in  a 
test-tube  by  displacement  of  air;  and  determine  if  it  is  inflammable,  and  if  it 
supports  combustion. 

4.  Suggest  an  experiment  to  prove  that  carbon  dioxide  is  moderately  soluble  in 
water  at  room  temperature.    Try  the  experiment.     When  CO2  dissolves  in  water 
the  reaction  is 

CO2(gas)    +   H2O   =   H2CO3(in    solution). 

Prove  by  an  experiment  that  the  reverse  reaction  can  take  place.  Under  what 
conditions  is  there  an  equilibrium?  Note.  At  room  temperature  a  solution  in 
equilibrium  with  CO2  gas  at  1  atmosphere  pressure  contains  about  0.04  mols 
H2COo  in  1  liter.  Question.  How  is  this  equilibrium  altered  by  an  increase  of 
temperature  ? 

5.  We  shall  now   consider  what  the  principal  substances  are  in  a  saturated 
carbonic    acid    solution.     Experiment.      Prepare    a    small    quantity    of    nearly 
saturated  solution  of  CO2  and  test  it  with  the  indicators,  phenolphthalein.  litmus, 
methyl  orange,   and  methyl  violet;   for  comparison   repeat  indicator  tests  with 
water  and  with  an  acid  solution  of  known  H+  concentration.     What  conclusion 
can  you  draw  with   regard  to  the   concentration  of  H+  in  the  saturated   CO2 
solution?    Note.     The  concentration  of  H+  in  this  solution  has  been  determined 
by   other   methods  to   be  about  0.0001   N    (W~4N).     From  this   value   for  the 
concentration  of   H+,  the  concentration  of   OH"  in  the  same   solution   can   be 
calculated  to  be  10  ~10  N;  explain  how  this  calculation  is  made,  cf .  Hildebrand, 
page  180. 

6.  Experiment.     Test  the  action   of   carbonic   acid  on   Ca++  by  passing  CO2 
gas  into  a  dilute  solution  of  CaCl2.    What  conclusion  can  you  draw  with  regard 
to  the  concentration  of  CO3~~  in  carbonic  acid  solution?    The  concentration  of 
CO3"  in  the  0.04  M  H2CO3  solution  has  been  found  by  other  methods  to  be 
of  the  same  order  of  magnitude  as  that  of  OH~.     Questions.     Making  use  of  the 
fact  that  the  only  ions  in  this  solution  are  produced  by  the  ionization  of  H2CO3 
(and  H2O),  what  conclusion  can  you  draw  with  regard  to  the  extent  to  which 
each  of  the  reactions   (/)  and  (^),  Paragraph  1,  has  taken  place  in  a  solution 
of  carbonic  acid?     What  is  the  approximate  concentration  of  HCO3"  in  the  0.04 
M  H2CO3  solution?     List  the  substances  present  in  this  solution  (a)  at  moderate 
concentration,  (b)  at  small  concentration (  two  substances),  and  (c)  at  extremely 
small  concentration  (two  substances). 

7.  The  neutralization  of  carbonic  acid.    When  NaOH  solution  is  added  gradu- 
ally to  a  carbonic  acid  solution  the  neutralization  of  H2CO3  takes  place  in  steps 
corresponding  to  the  two  steps  in  the  dissociation  of  H2CO3.    The  main  reaction 
for  the  first  step  in  the  neutralization  is 

HoCOs  +  OH-  +  (Na+)  =±*  H2O  +  HCCV  +  (Na+), 

and  the  solution  made  from  eaual  molal  quantities  of  H2CO3  and  NaOH  is  a 

sodium  bicarbonate  soultion.     The  second  step  in  the  neutralization  of  H2CO3  is 

HCO3-  +  OH-  ==  H2O  +  CO3-  - 

[39] 


This  is  the  reaction  for  the  neutralization  of  HCO3~  by  a  strong  base,  and  was 
considered  in  Assignment  25. 

8.  Substances  present  in  a  solution  of  sodium  carbonate.    Review  your  notes, 
Assignment  25,  on  the  relative  concentrations  of  CO3~  ~,  HCO3~  and  OH~  in  a 
0.5  M  Na2CO3  solution;    and  calculate  the  concentration  of  H+  in  this  solution. 
The  concentration  of  H2CO3  in  this  solution  is  also  extremely  small,  but  of  course 
is  greater  than  the  concentration  of  H+.    Questions.     How  would  you  prepare  1 
mol  of  solid  Na2CO3  from  1  mol  of  NaHCO3,  and  also  from  1  mol  of  H2CO3? 
What  quantity  of  NaOH  would  be  required  in  each  case?     List  the  substances 
which  are  present  in  a  solution  of  Na2CO3,  (a)  at  moderate  (or  high)  concentra- 
tion, (b)  at  relatively  small  concentration  (two  substances),  and  (c)  at  extremely 
small  concentration   (two  substances). 

9.  Substances  present  in  a  solution  of  sodium  bicarbonate.    From  the  fact  that 
HCO3~  is  a  weak  acid,  what  would  you  expect  to  be  the  principal  substances  in 
a  solution  of  NaHCO3?     Noting  that  the  substance  HCO3~  is  intermediate  in 
composition  between  H2CO3  and  CO3~~,  and  considering  the  equilibria  involved, 
would  you  expect  the  concentration  of  CO3~"  in  a  solution  of  NaHCO3  to  be 
greater  or  less  than  in  a  solution  of  (a)  H2CO3,   (b)   Na,CO3?     The  following 
experiment,  on  the  precipitation  of  CaCO3  by  NaHCOs  solution  at  room  temp- 
erature, furnishes  evidence  with   regard  to  the  concentration  of  CO3~~  in  this 
solution,    (but  it  should  be   remembered  that  more  CaCO3  is  precipitated  than 
corresponds  to  the  actual  concentration  of  the  CO3~~  since  on  account  of  the 
displacement  of  equilibrium  some  more  CO3~~  forms  as  the  CaCO3  is  precipitated. 
Experiment.     To  10  cc.  M  NaHCO3  in  a  flask  add  about  50  cc.  water,  and  10  cc. 
N  CaCl2 ;  shake  the  mixture.     Filter  and  note  the  amount  of  the  CaCO3  preci- 
pitate.   Compare  the  result  with  that  obtained  in  the  experiment  with  Ca++  and 
H,CO3,  Paragraph  6.    Repeat  the  experiment  using  10  cc.  TV  (0.5  M)   Na2CO3 
instead  of  10  cc  M  NaHCO3.     Heat  the  two  nitrates  to  boiling,  and  test  the 
gas  evolved  in  the  NaHCO3  experiment  by  passing  it  into  a  clear  solution  of 
Ca(OH)2    or    Ba(OH)2    (cf.    Assignment    32,    Paragraph    11).      Compare   the 
amount  of  the  precipitate  now  obtained  with  that  formed  in  the  cold  solution. 
Complete  the  following  equation : 

Ca++  +  2HCO3-  =  CaCO3  (solid) 

and  suggest  an  explanation  for  the  shifting  of  the  equilibrium  when  the  mixture 
is  heated  to  boiling.  The  reaction  that  takes  place  when  a  solution  of  NaHCO3 
is  heated  to  boiling  is 

2HC03-  =  H2C03  +  C03"  =  C02  (gas)  +  H2O  +  CO3" 

Questions.  How  will  the  concentrations  of  H2CO3  and  CO3~~  in  a  NaHCO3 
solution  be  changed  (/)  when  the  solution  is  boiled,  and  (2)  when  CaCL 
solution  is  added  at  room  temperature?  What  conclusions  can  you  draw  with 
regard  to  the  relative  concentrations  of  HCO3~,  and  H2CO3  and  CO3"  ~  in  a  freshly 
prepared,  cold  solution  of  NaHCCX? 

10.  Experiment.    Test   a   freshly   prepared,    approximately   molal    solution   of 
NaHCO3  with  litmus  and  with  phenolphthalein  and  estimate  approximately  the 
concentration  of  OH~  and  H+  in  the  solution.     Repeat  the  experiment  with  a 
solution  which  has  been  heated  to  boiling,  or  which  has  been  allowed  to  stand 
in  the  laboratory  for  several  days.     From  your  results  in  Paragraphs  9  and  10 
state  which  substances  are  present  in  a  pure  NaHCO..-  solution,  (a)  at  moderate 
(or  high)  concentration,  (ft)  at  relatively  small  concentration,  (two  substances) 
and  (c)  at  still  smaller  concentration  (two  substances). 

11.  Summarize  in  tabular  form  the  lists  referred  to  in  the  last   sentence  of 
each  of  the  three  Paragraphs,  6,  10  and  8,  arranging-  the  columns  in  the  order 
H,CO3,  NaHCO3  and  Na2CO3.     Note  the  regular  change  in  the  concentration 
of  "each  substance,  e.  g.  H2CO3,  when  the  three  solutions  are  considered  in  this 

[40] 


order.  If  the  arrangement  of  the  various  substances  in  your  table  is  not  sym- 
metrical you  have  probably  made  some  mistake.  Question.  What  is  the  equation 
for  the  main  reaction  when  0.1  mol  of  strong  acid  is  added  to  (a)  0.1  mol 
Na,CO3,  and  (b)  0.05  mol  Na2CO3? 

12.  The  decomposition  of  solid  sodium  bicarbonate.    Experiment.    Place  about 
1   gram  NaHCO3   in  a  hard-glass   test-tube  and  lead  the  delivery  tube  into   a 
solution  of  Ba(OH)2,  or  of  CaCL  and  NH4OH.     Heat  the  tube  until  no  more 
gas  is  given  off,  remove  the  delivery  tube  from  the  solution,  and  allow  the  hard- 
glass  tube  to  cool.     What  is  the  evidence  that  CO2  was  formed?     That  water 
was  formed  ?    What  must  the  residue  be  if  H2O  and  CO2  are  formed  in  equimolal 
quantities?     (This  experiment  may  be  performed  quantitatively  by  weighing  the 
hard-glass  tube  (/)  empty,  (2)  with  NaHCO3,  and  (5)  with  the  residue.) 

Experiment.  Continued.  Clean  the  delivery  tube  and  lead  it  into  a  fresh 
solution  of  Ba(OH)2.  Pour  water  carefully  into  the  hard-glass  tube  until  it  is 
half  full,  but  do  not  stir  the  mixture.  Add  2  cc.  6  N  HC1,  and  at  once  insert  the 
stopper.  What  is  the  gas  evolved?  Warm  the  mixture  slowly  by  placing  the 
hard-glass  tube  in  a  beaker  of  water  and  heating  the  latter.  Questions.  If  the  gas 
volumes  had  been  measured,  what  would  have  been  the  ratio  of  the  volumes  in 
the  two  parts  of  the  experiment?  If  all  the  CO2  evolved  had  been  converted  into 
BaCO3,  what  would  have  been  the  relative  amounts  obtained  in  the  two  parts 
of  the  experiment? 

13.  Reaction  between  an  acid  and  the  negative  ion  of  a  weaker  acid.     Give 
examples  of  the  action  of  a  strong  acid  on  a  solution  of  the  salt  of  a  weak  acid. 
Experiment.        Determine     the     action    of    dilute    acetic    acid    on     NaHCO3. 
Which  is  the  stronger  acid,  HAc  or  H2CO3  ?    Taking  into  consideration  the  fact 
that  HCO3~,  the  negative  ion  of  H2CO3,  is  also  a  weak  acid  (whose  ions  are  H* 
and  CO3— ),  and  the  additional  fact  that  H2CO3  is  a  stronger  acid  than  HCO3', 
predict  what  will  happen  when  H2CO3  is  introduced  into  a  solution  of  a  car- 
bonate. 

14.  Experiment.     Test  your  answer  to  the  last  question  by  placing  a  small 
amount  of  freshly  precipitated  CaCO3  in  about  100  cc.  water  and  saturating  the 
mixture  with  CO2.     Filter  off  any  CaCO3  that  remains  and  heat  the  filtrate  to 
boiling. 

15.  Summarize  all  the  evidence,  presented  in  this  Assignment  that  the  reaction 

2HC03-  =  H2C03  +  C03" 

is  reversible,  noting  especially  the  experimental  conditions  under  which  the  re- 
action will  proceed  almost  completely,  (a)  as  written,  and  (b)  in  the  reverse 
direction.  Question.  Which  solution  would  be  more  alkaline  at  room  tempera- 
ture (on  account  of  hydrolysis),  a  molal  solution  of  a  salt  of  a  weak  monobasic 
acid  of  the  same  strength  as  H2CO3,  or  a  molal  solution  of  NaHCO3  ?  Give  your 
reasoning. 

16.  Problems.  (/)     Can  the  following  substances  be  present  at  moderate  con- 
centrations in  the  same  solution?     If  not,  what  is  formed? 

(a)  H2CO3   and   Ca++  (/)   HCO3-  and  H+ 

(b)  H2CO3  and  OH'  (g)  HCO3-  and  OH- 

(c)  H2CO3  and  CO3"  (h)  HCCV  and  CO," 
(<0   H2CO3  and  HCO,-  (i)  CO8—  and  OH- 
(e)   H,CCL  and  HH 

(2)  What  weight  of  Na2CO3  can  be  obtained  from  8.4  g.  NaHCO3,  (a)  by 
heating,  and  (b)  by  treating  the  bicarbonate  with  NaOH  solution?  What  is 
the  least  volume  of  N  NaOH  that  can  be  used  in  (b)  ? 

[41]      . 


(j)  What  volume  of  CO2  at  standard  conditions  can  be  obtained  from  8.4  g. 
NaHCO3,  (a)  by  heating,  and  (b)  by  treating  the  bicarbonate  with  HC1  solution? 
What  is  the  least  volume  of  N  HC1  that  can  be  used  in  (b)  ? 

(4)  When  a  current  of  CO2  is  passed  into  a  solution  of  Ca(OH)2  a  precipitate 
is  observed  to  form,  and  then  dissolve.  If  the  final  solution  is  now  heated  to 
boiling,  a  precipitate  again  appears.  State  what  has  happened  in  this  experiment, 
and  write  an  equation  for  the  main  reaction  in  each  of  these  stages. 


ASSIGNMENT  34 
SULPHATES,  CHLORIDES  AND  NITRATES  OF  COPPER  AND  ZINC 


1.  In  Assignment  34  and  the  four  following  Assignments  we  shall  study  the 
chemistry  of   copper  ion,   Cu++,  silver  ion,  Ag+,  and  zinc  ion,  Zn++.     We  shall 
first  devote  our  attention  to  the  common  soluble  salts,  and  to  the  method  of 
transforming  one  salt  into  another ;  and  then  we  shall  study  the  sparingly  soluble 
compounds,  the  methods  of  dissolving  them,  and  the  equilibria  involved.     The 
same  method  of  treatment  will  be  used  later  in  studying  the  chemistry  of  ions 
of  other  metals,  and  many  of  the  results  now  obtained  in  studying  copper,  silver 
and  zinc  are  also  true  for  other  metals.    Read  again  Paragraph  2,  Assignment  31. 

2.  Solubility    of    nitrates,    acetates,    chlorides    and   sulfates    of   metals.      The 
nitrates  and  acetates  of  all  metals  are  soluble  in  water  and  this  is  also  true  foi 
nearly  all  the  chlorides  and  sulfates.     Question.     Give  an  example  of  a  diffi- 
cultly soluble  chloride  and  of  a  difficultly  soluble  sulfate.     Give  the  formulas  of 
the  nitrates,  acetates,  chlorides  and  sulfates  of  copper,  silver  and  zinc. 

3.  Preparation   of  sulfates  from   nitrates  and   chlorides.     Experiment.     Test 
the  relative  volatility  of  HC1,  HNO3  and  H2SO4  by  evaporating  a  few  drops  of 
a  concentrated  solution  of  each  of  these  acid's  in  a  casserole  (out  of  doors  or  in 
a  fume  closet).     Cf.  Hildebrand,  pages  173-175.     Suggest  a  method  of  preparing 
solid   copper  sulfate  from  copper  nitrate  based  on   the  difference  in  volatility 
of  HNO3  and  H2SO4.  Experiment.    Try  your  method  with  15  cc.  of  the  labora- 
tory solution  of  copper  nitrate.     Recrystallize  the  copper  sulfate  by  mixing  the 
residue,  after  it  is  cold,  with  2  or  3  cc.  water,  heating  the  mixture  to  boiling 
and  letting  it  cool  slowly.     Wash  the  crystals  with  a  very  little  water,  test  a 
small  portion  for  NO3",  and  save  the  remainder. 

4.  Suggest  a   similar  method   for  the  preparation   of  solid  zinc   sulfate   from 
zinc  chloride.     How  would  you  test  whether   the  final  product   is   free   from 
chloride  ? 

5.  The   problem   of   preparing  a   soluble   chloride   or   nitrate    from   a   soluble 
sulfate  will  be  considered  in  the  following  Assignment. 

6.  Conversion  of  soluble  chlorides  into  nitrates,  and  nitrates  into   chlorides. 
Experiment.    Mix  2   cc.   concentrated   HNO3   solution  and   6   cc.   concentrated 
HC1  solution,  and  let  the  mixture  stand.     Evaporate  a  few  drops  of  the  solution 
to  dryness   (out  of  doors  or  in  a  fume  closet)   and  note  if  there  is  a  residue. 
The  gradual  deepening  of  the  color  of  the  solution  at  room  temperature  proves 
not  only  that  a  reaction  is  taking  place  but  that  this  reaction  is  not  a  rapid  one. 
Questions.     How  would  the  speed  of  the  reaction  be  altered  (a)  by  raising  the 
temperature,  and   (b)   by  using  less  concentrated  acids?     Note.     Although  this 
reaction  does  not  take  place  in  a  dilute   solution  which   contains  H+,   Cl~  and 
NO3",  it  can  be  made  to  do  so  by  concentrating  the  solution  by  evaporation  to 
a  small  volume.     Since  both  Cl~,  and  NO3~  are  used  up  in  this  reaction  and  the 
products  are  volatile,  we  can  make  use  of  this  reaction  (/)  to  remove  Cl~  from 
a  solution  by  heating  with  excess  of  concentrated   HNO3,   or    (2)    to   remove 

[42] 


NO3~  from  a  solution  by  heating  with  excess  of  concentrated  HC1.  Question. 
If  excess  of  concentrated  HNO3  is  added  to  a  small  amount  of  NaCl  and  the 
mixture  evaporated  just  to  dryness,  what  would  you  expect  the  solid  residue 
to  be.  Note.  The  mixture  of  concentrated  HNO3  and  HC1  is  called  aqua  regia, 
and  the  reaction  that  takes  place  is  NO8-  +  3  Cl~  +  4H+  =  NOC1  +  CL,  +  H2O. 

7.  Experiment.     To  about  0.2  g.  copper  chloride  (or  2  cc,  N  ZnCl2  solution) 
in  a  casserole,  add  3  cc.  concentrated  HNO3  and  stir;  add  one  or  two  drops  of 
this  mixture  to  10  cc.  water,  test  this  solution  for  Cl~  and  set  it  aside  for  com- 
parison  with   later   results.      Evaporate   the   mixture   of   the   chloride   and   con- 
centrated HNO3  just  to  dryness  (out  of  doors  or  in  a  fume  closet).,  add  5  cc. 
concentrated    HNO3    being    careful    to    dissolve    any    solid   on    the    side   of   the 
dish,  and  test   one  or  two    drops   of    the    solution    for    Cl~   as   before.       Again 
evaporate  the  mixture  just  to  dryness,   and  test  a  portion  of  the  residue  for 
chloride.     Questions.     What  is  the  final   solid   residue?     What  conclusion  can 
your  draw  from  a  comparison  of  the  amounts  of  the  AgCl  precipitates  in  your 
three  tests?     Does   this  experiment   furnish   any   evidence  that   the   reaction  is 
slow,  even  at  a  temperature  in  the  neighborhood  of  100°,  when  the  concentration 
of  one  of  the  reacting  substances  is  small? 

8.  How  would  you  convert  copper  nitrate,  or  zinc  nitrate,  into  the  chloride? 

9.  Treatment  of  silver  chloride  with  hot  concentrated    nitric    acid    and    sul- 
furic  acid.    Experiment.    Prepare  some  silver  chloride,  and  wash  it  with  dilute 
HC1  and  then  with  water.     Place  a  small  amount  of  the  silver  chloride  in  a 
porcelain  dish,  and  3  or  4  cc.   concentrated  HNO3,  evaporate  the  mixture  to 
dryness.     Add  a  little  water  and  test  the  solution  for  Ag+.     Repeat  the  experi- 
ment  with    the   residue.      Compare   the   result   with   that   in   the   experiment   in 
Paragraph  7,  and  state  how  the  speed  of  the  reaction  between  Cl~  and  NO3~  in 
acid  solution  depends  on  the  concentration  of  Cl~. 

10.  Test  the  action  of  hot  concentrated  H2SO4  on  silver  chloride  by  heating 
for  several  minutes  a  small  amount  of  the  salt  with  3  or  4  cc.  of  the  acid  in  a 
porcelain  dish  covered  with  a  watch  glass.     Set  the  dish  aside  until  it  is  cool, 
pour  the  mixture  into  water,  and  test  the  solution  for  Ag+.    Note.     The  chloride 
is  expelled  as  HC1  gas  in  this  experiment,  but  since  the  temperature  may  have 
exceeded  300°  it  is  not  surprising  that  the  result  is  different  from  that  obtained 
in  the  experiment  with  AgCl  and  concentrated  HNO3. 

11.  Problems.    (/)   What  will  be  the  solid  residue  when  each  of  the  follow- 
ing solutions  is  evaporated  to  dryness? 

(a)  A  solution  which  contains  Ag+,  H+  and  NO3~. 

(b)  A  solution  which  contains  Cu++  and  NO3~  at  small  concentrations  and 
H+  and  SO4'~  at  large  concentrations. 

(c)  A  solution  which  contains  Cu++  and  NO3~  at  large  concentrations  and 
H+  and  SO4~~  at  small  concentrations. 

(d)  A  solution  of  zinc  sulfate  to  which  nitric  acid  has  been  added. 

(2)  How  would  you  prepare  solid  silver  sulfate  from  (a)  solid  silver  nitrate, 
(b)  solid  silver  acetate? 

(j)  When  solid  copper  sulfate  is  prepared  by  crystallization  from  an  aaueous 
solution  it  contains  water  of  crystallisation  and  its  formula  is  CuSO4  •  5H2O. 
When  this  blue  substance  is  heated,  white  anhydrous  CuSO4  is  finally  obtained. 
Write  the  equation.  The  reverse  reaction  takes  place  when  air  saturated  with 
water  vapor  is  passed  over  the  anhydrous  salt  at  room  temperature.  If  a  vessel 
containing  CuSO4  •  5H2O  were  evacuated  for  a  few -minutes  and  then  closed, 
what  substances  would  be  present  in  the  closed  vessel?  What  reason  do  you 
have  for  the  conclusion  that  the  pressure  of  water  vapor  in  the  vessel  at  room 
temperature  must  be  less  than  the  partial  pressure  of  water  vapor  .at  the  same 

F431 


temperature?  How  would  the  pressure  in  the  vessel  be  altered  by  raising  the 
temperature?  Give  two  other  examples  of  a  reversible  reaction  in  which  one 
solid  substance  dissociates  into  another  solid  and  a  gas,  and  in  each  case  give 
an  approximate  value  for  the  equilibrium  pressure  at  some  definite  temperature. 
(Refer  to  the  lectures  on  the  dissociation  of  calcium  carbonate  and  calcium 
hydroxide  when  heated.) 


ASSIGNMENT  35 
HYDROXIDES  OF  COPPER,  SILVER  AND  ZINC 


References.     Hildebrand,   Chapter  V  pages   72-73.  80-82;   Chapter   XIII  pages 

194-195. 


1.  We   have    learned    from    the   laboratory   work   and    the   lectures   that,   the 
hydroxides  of  sodium,   potassium  and  the  other  alkali  metals,   and  ammonium 
hydroxide  are  readily  soluble  in  water,  and    that    the    hydroxides    of    barium, 
strontium   and   calcium   are   moderately  soluble    (with  the   solubility   decreasing 
rather  rapidly  in  the  order  named).     The  hydroxides  of  all  other  metals  are 
difficulty  soluble  in  water. 

2.  Cuprlc    hydroxide.      State    what    solutions    you    would   mix    to    prepare 
Cu(OH)2.     Point  out  the  substances  involved  in  the  equilibrium  between  the 
solid  and  its  saturated  solution  and  predict  what  will  happen  when  the  solid  is 
treated    with    HC1,    HNO3    or    H2SO4    solution.      Experiment.       Test    you* 
answer  by  preparing  some  copper  hydi  oxide  and  treating  it  with  acids.     Also 
determine  if  it  dissolves  in  0.5  'N  NaOH. 

3.  Cuprlc  oxide.       Experiment.       Collect    some    Cu(OH)2    on    a    filter    and 
wash  it  once  with  water.     Mix  some  of  the  Cu(OH)2  with  water  and  heat  the 
mixture  to  boiling.     Place  the  remainder  of  the  Cu(OH)2  in  a  procelain  dish 
and  heat  gently.    Cupric  oxide  has  formed.     Does  the  reverse  reaction  between 
CuO  and  H2O  take  place  at  room  temperature?     Predict  what  will  happen  when 
NaOH  solution  is  added  to  a  solution  of  Cu(NO3)2  at   100°,  and  test  your 
answer.    Test  the  action  of  HC1,  HNO3  or  H2SO4  solution  on  CuO.     Question. 
How  would  you  prepare  solid  CuCl2  from  Cu(OH)2  or  CuO? 

4.  Sliver    oxide.        Experiment.        To    10    cc.    0.1    TV    AgNO3    solution    add 
NaOH  solution  slowly  until,  after  shaking,  the  solution  reacts  strongly  alkaline 
to   litmus.      Silver   oxide,   Ag2O,   is    formed.      Filter,    acidify   the    filtrate   with 
HNO3  and  test  it  for  Ag+.     Write  the  equation  for  the  reaction  between  Ag+ 
and  OH~  to  form  Ag2O.     Point  out  the  substances  involved  in  the  equilibrium 
between  the  solid  substance  and  its  saturated  solution,  and  predict  what  will 
happen  when   the  solid  substance  is  treated  with  HNO3  solution.     Test  your 
answer  by  an  experiment. 

5.  Relation  between  oxides  and  hydroxides  of  metals.     Contrast  the  action  of 
water  or  water  vapor  at  room  temperature  on  calcium  oxide,  and  on  copper 
oxide  or  silver  oxide.     Review  Problem  3,  Assignment  34;  and  note  that  the 
reaction 

Oxide  of  a  metal  (solid)  +  H2O  (gas)  =  Hydroxide  of  a  metal  (solid) 
takes  place  completely  in  the  case  of  oxides  of  the  alkali  metals,  as  Na2O,  does 
not  have  any  tendency  to  take  place  in  the  case  of  oxides  of  the  more  noble 
metals,  as  Ag2O,  and  can  be  shown  to  be  reversible  for  many  oxides  which 
are  intermediate  between  these  two  extremes.  (In  many  cases  equilibrium  is 
reached  only  very  slowly,'  especially  at  low  temperatures.)  On  account  of  this 
relation  between  oxides  and  hydroxides  of  metals,  it  is  not  surprising  that  oxides 
of  metals  in  general  are  capable  of  neutralizing  acids.  Questions.  What  is 
anhydride  of  a  base?  Give  examples.  Is  it  always  correct  to  say  than  an 

[441 


anhydride  of  a  base  will  react  with  water  to  form  the  base?  What  is  an  acid 
anhydride!'  Give  examples.  Also  give  examples  of  reactions  between  an  acid 
anhydride  and  a  base,  and  between  an  acid  anhydride  and  a  basic  anhydride. 

6.  Acids  which  contain  oxygen  may  be  considered  to  be  related  to  the  hydrox- 
ides of  non-metals.  Thus  H2SO4  may  be  regarded  as  H6SO6  or  S(OH)0  from 
which  2  molecules  of  water  have  been  withdrawn 

S(OH)6  =  H2SO4  +  2  H2O; 

and  the  formula  of  phosphorous  acid,  H3PO3,  may  be  written  P(OH)3.  The 
characteristic  properties  of  these  acids,  to  yield  negative  ions  containing  oxygen 
on  dissociation  or  neutralization,  is  then  to  be  attributed  to  the  existence  of 
strong  bonds  between  the  non-metal  and  the  oxygen  in  the  compound;  which  is 
really  due  to  the  great  tendency  of  non-metals  to  hold  electrons  firmly.  On  the 
other  hand  metals  in  compounds  do  not  hold  electrons  firmly,  and  positive  ions 
of  metals  are  formed  when  bases  or  salts  dissociate.  However,  certain  elements 
which  form  a  positive  ion  also  can  form  a  negative  ion  containing  oxygen;  the 
hydroxide  of  such  an  element  has  the  properties  of  a  base  in  that  it  can  neutralize 
an  acid,  and  also  has  the  properties  of  an  acid  since  it  can  neutralize  a  base:  it 
is  said  to  be  amphoteric. 

7.  Zinc    hydroxide.       Experiment.       Prepare    some    Zn(OH)2    by  treating  a 
solution    containing   Zn++   with   a   very   small   amount   of   NaOH.     Collect   the 
Zn(OH)2  on  a  filter,  treat  a  portion  with  a  strong  acid,  and  another  portion  with 
NaOH  or  KOH  solution.        Question.        What  evidence  is   furnished  by  this 
experiment  that  zinc  hydroxide  is  an  amphoteric  substance?     The  reaction  be- 
tween zinc  hydroxide  and  excess  OH~  is  analogous  to  the  first  stage  in  the  neu- 
tralization of  carbonic  acid : 

H0ZnOo  (solid)  +  OH-  =  H,O  +  HZnCX- 
H2COS  +  OH-  —  H2O  +  HCOS- 

The  negative  ion,  HZnO3",  is  usually  zincate  ion,  though  strictly  speaking  this 
name  should  be  reserved  for  the  ion  ZnO2~~.  The  latter  substance  probably 
forms  to  some  extent  when  the  concentration  of  OH"  is  very  large.  Solid 
NaHZnO,  has  not  been  prepared,  but  Na^ZnOo  can  be  made  by  fusing  zinc  oxide 
with  NaOH. 

8.  Predict  what  will  happen  when  a  solution  of  a  strong  acid  is  added  slowly 
to  sodium  zincate  solution.     Test  your  answer  by  an  experiment. 

9.  Preparation  of  a  nitrate  or  chloride  from  a  soluble  sulfate.       Experiment. 
Prepare  a  solution  of  copper  nitrate  from  copper  sulfate  solution  by  precipitat- 
ing copper  hydroxide  from  a  very  dilute  CuSO4  solution,  washing  the  Cu(OH)2 
precipitate,  and  transforming  it  into  the  nitrate.     Test  the  product  for  sulfate. 

10.  Suggest  a   second   method   based   on  the   removal  of   the   sulfate  ion  by 
means  of  barium  ion   (cf.  Hildebrand.  pages  207-208).     Try  your  method,  and 
test  the  copper  nitrate  solution  for   SO4~~  and  for  Ba++.     Which  is  the  better 
method  to  use  on  a  small  scale  in  the  laboratory? 

11.  Suggest  two  methods  of  preparing  solid  zinc  chloride  from  zinc  sulfate. 

12.  Problems.   (/)     Write  equations  for  the  action  of  HC1  solution  in  excess 
on  calcium  oxide,  on  ferrous  and  ferric  oxides  (FeO  and  Fe.X)3),  on  ferrous  and 
ferric  hydroxides,  and  on  sodium  zincate  solution. 

(2)  A  solution  containing  0.2  mol  Na2SO4  is  mixed  writh  a  solution  contain- 
ing 0.1  mols  Ba(NO3)2,  the  precipitate  is  removed  by  filtration,  the  wash  water 
is  run  into  the  filtrate  until  the  final  volume  of  the  latter  is  1  liter.  What  sub- 
stances are  present  in  the  final  solution  and  what  is  the  concentration  of  each? 
Suggest  a  method  of  preparing  pure  sodium  nitrate  from  sodium  sulfate. 

( j)    Give  two  methods  of  preparing  CuCL  from  CuSO4. 

[45] 


(4)   Give  an  example  of  an  amphoteric  hydroxide  other  than  Zn(OH)2,  and 
write  the  equation  for  its  reaction  with  (a)  a  strong  acid,  (b)  a  strong  base. 


ASSIGNMENT  36 
COMPLEX  IONS  OF  COPPER,  SILVER  AND  ZINC  WITH  AMMONIA 


References.    Hildebrand,  pages  184  and  202-205. 

1.  The  ions  of  certain  metals   have  the  power  of   forming  compounds   with 
NH3,  which   are  examples  of  "complex  ions".     The  ammonia  is   supplied  by 
adding   NH4OH   solution.     There   is  an  equilibrium  between   the   complex  ion, 
NH3  or  NH4OH,  and  the  ion  of  the  metal;  and  when  the  NH4OH  is  present 
in  excess  the  concentration  of  the  ion  of  the  metal  is  often  very  small.     Some 
examples  will  be  considered  in  the  present  Assignment. 

2.  Experiment.     Dilute  5   cc.  N  CuSO4  with  20  cc.  water  and  add  two  or 
three  drops  6  N  NH4OH.     The  precipitate   is   Cu(OH)2.      Question.      What 
evidence  with  regard  to  the  solubility  of  Cu(OH)2  in  water  is  furnished  by  this 
experiment?     (Consider  the  low  concentration  of  OH~  in  NH4OH  solution  and 
the  result  of  the  experiment  in  Paragraph  6,  Assignment  32. 

3.  Continue  to  add  the  NH4OH  solution  until,  after  shaking,  the  solution  is 
clear.    Observe  the  volume  of  NH4OH  used.    What  is  the  evidence  that  a  new 
substance  is   formed?     Repeat  the  experiment  with  the  same  volume  oi  6  N 
NaOH  instead  oi  6  N  NH4OH.     State  the  reasoning  by  which  you  may  con- 
clude that  the  substance  obtained  in  the  experiment  with  NH4OH  solution  is 
not  formed  by  a  reaction  between  OH~  and  Cu(OH)2  similar  to  that  between 
OH~  and   Zn(OH)2.     Try   an   experiment   with   CuSO4    solution    and    NH4C1 
in  order  to  determine  if  Cu++  reacts  with  NH4+  in  the  same  way  as  with  NH3 
or  NH4OH. 

4.  The  new  substance  formed  when  Cu(OH)2  dissolves  in  NH4OH  solution 
is  Cu(NH3)4++.     The  final  solution  contains  NH4OH,  and  OH-  must  be  present 
at  an  appreciable  concentration ;  it  may  be  concluded  that  the  concentration  of 
Cu++  is  extremely  small,  since,  otherwise,  Cu(OH)2  would  precipitate.  Questions. 
Which  solution  has  the  greater  concentration  of  Cu++,  a  saturated  solution  of 
Cu(OH)2  or  a  solution  of  the  complex  ion  containing  NH4OH?     How  is  the 
equilibrium  between  Cu(OH)2  and  its  saturated  solution  affected  when  NH.,OH 
is  added? 

5.  The  equation   for  the   main   reaction  when  a  solution  of  a   copper  salt  is 
treated  with  excess  NH4OH  is 

Cu++  +4NH4OH  =  Cu(NH3)4++  +  4H2O. 

A  simple  method  of  demonstrating  that  the  reverse  reaction  will  take  place  is  to 
remove  the  NH4OH.  Suggest  two  methods  of  doing  this  and  try  your  methods. 

6.  Zinc  and  silver  also  form  complex  ions  with  ammonia, 

Zn(NH8)4"  and  Ag(NH3)2+. 

Experiment.  To  a  dilute  solution  of  ZnSO4  add  NH4OH  solution  drop  by 
drop,  shaking  the  mixture  after  each  addition  of  NH4OH;  when  some  but  not 
all  of  the  Zn(OH)2  has  dissolved  test  the  solution  with  trinitrobenzol  to  determine 
approximately  the  OH~  concentration.  Repeat  the  experiment,  using  NaOH 
instead  of  NH4OH,  and  note  again  the  OH-  concentration  in  the  solution  when 
Zn(OH)2  is  partly  but  not  completely  dissolved.  Question.  From  a  comparison 
of  the  OH~  concentration  at  equilibrium  in  the  two  cases  what  conclusion  can 
you  draw  with  regard  to  the  presence  of  zinc  as  a  different  substance  in  the 
NH4OH  and  NaOH  solutions? 

[46] 


7.  Experiment.      Prepare    some    silver   oxide   and    some    silver    chloride,    and 
treat  each  substance  with  NH4OH  solution.     Predict  what  will  happen  when 
the  resulting  solutions  are  acidified  with  nitric  acid,  and  test  your  answers  by 
experiments. 

8.  Which  solution  would  contain  a  greater  concentration  of  Ag+,  a  saturated 
solution  of  silver  acetate  (see  Assignment  24,  Paragraph  5)  or  a  solution  which 
contains  the  complex  ion  and  NH4OH?     What    then    must    be    the    action    of 
NH4OH  solution  on  solid  AgAc? 

9.  Questions.    When  AgCl  is  dissolved  in  excess  NH4OH  solution,  what  is  the 
principal  positive  ion,  and  the  principal  negative  ion,  in  the  resulting  solution? 
When  o.i  mols  AgCl  are  dissolved  and  the  final  volume  is  100  cc.  what  is  the 
approximate  concentration  of  each  of  these  ions?     Is  the  concentration  of  Ag* 
smaller  or  larger  than  in  a  saturated  solution  of  AgCl  in  water? 

10.  When   Cu(OH)2   is   dissolved  in   excess   NH4OH   solution  what   is  the 
principal  negative   ion  in   the   final   solution?   ,  It   is   important   to   realize   that 
Cu(NH3)4(OH)2,    Zn(NH3)4(OH)2    and   Ag(NH3)2OH    are    soluble,    strong 
bases. 

11.  We  shall  next  consider  the  action  of  NH4OH  solution  on  silver  iodide,  a 
salt  which  is  much  less  soluble  in  water  than  either  silver  oxide  or  silver  chloride. 
Solid  Agl,  suspended    in  water,    is  in    equilibrium    with    Ag+    and    I~    in    the 
saturated  solution.     When  NH4OH  is  added  some  Ag+  must  react  with  it  to 
form  Ag(NH3)2+,   and  some  solid  Agl   must  dissolve   to   establish   again   the 
solubility  equilibrium.    We  can  therefore  predict  that  Agl  must  be  more  soluble 
in  NH4OH  solution  than  in  water,  but  it  is  not  safe  to  predict  that  a  large 
amount  of  it  will  dissolve.     Experiment.     Prepare  a  solution  of  silver  oxide 
or  silver  chloride  in  NH4OH  and  add  to  this  solution  a  few  drops  of  0.1  N  potas- 
sium iodide  solution.     In  another  experiment  collect  some  solid  silver  iodide  on 
a  filter  paper,  wash  it  with  water,  and  treat  it  with  dilute  NH4OH  solution. 
Devise  and  try  an  experiment  with  the  solution  thus  prepared  to  determine 
whether  an  appreciable  amount  of  Agl  has  dissolved. 

12.  Predict  what  will  happen  when  silver  chloride  is  treated  with  a  solution 
of  potassium  iodide.     Try  the  experiment. 

13.  The  ions  of  the  alkali  metals,  as  Na+,  K+,  etc.;  those  of  the  alkaline  earth 
metals,  as  Ca++,  SrTT  and  Ba++;  and  many  other  positive  ions  do  not  combine 
with  NH3  to  form  complex  ions.     Question.    Will  a  difficultly  soluble  salt  of 
calcium,  as  CaCO,  or  CaC2O4,  dissolve  in  NH4OH  solution  to  a  greater  extent 
than  in  water? 

14.  Problems,     (i)     Write  equations  for  the  main  reactions  in  the  following 
experiments : 

(a)  A   few   drops  of   NH4OH   solution  are  added   to  a   solution   contain- 
ing Zn++. 

(b)  Excess  NH4OH  solution  is  added  to  a  solution  containing  Zn++. 

(c)  An  ammoniacal  solution  of  silver  chloride  is  acidified  with  nitric  acid. 

(2)  The  solubilities  of  AgCl  and  Agl  in  water  are  10~4  and  10~7  mols  per 
liter,  respectively.  In  two  experiments  with  NH4OH  solution  each  substance 
is  found  to  be  one  thousand  times  more  soluble  than  in  water.  How  many  grams 
of  each  has  dissolved  in  i  liter  of  the  NH4OH  solutions? 

(j)   Give  two  examples  of  complex  ions  other  than  those  with  ammonia. 

(4)  From  the  positions  of  Cu,  Ag,  and  Zn  in  the  Periodic  System  (Hildebrand 
page  257),  state- what  other  ions  might  be  expected  to  form  ammonia  complexes. 
Check  your  answer  by  referring  to  Hildebrand,  page  269. 

[47] 


ASSIGNMENT  37 
CARBONATES  AND  SULFIDES  OF  COPPER,  SILVER  AND  ZINC 


References.     Hildebrand,  pages  199-201  and  195-197. 

1.  In    Assignment    33    we    learned    that    when    carbon    dioxide    dissolves    in 
water  the  solution  contains  the  weak  acid  H2CO3,  and  in  the  saturated  solution 
we  recognized  an  equilibrium  between  H2CO3  in  the  solution  and  CO2  in  the 
gas  space.     We  found  that  carbonic  acid  ionizes  in  two  steps  and  can  be  neu- 
tralized in  stages  to  give  solutions  of  the  two  corresponding  types  of  salts,  and 
that  when  these  salts  are  treated  with  a  strong  acid  carbon  dioxide  is  evolved. 
In  the  present  Assignment  we  shall  continue  the  study  of  difficultly  soluble  car- 
bonates, and  we  shall  deal  with  the  salts  of  another  dibasic  acid,  H2S,  which  is 
in  many  respects  analogous  to  carbonic  acid. 

2.  Hydrogen  sulfide,  like  carbon  dioxide,  is  a  gas  which  is  soluble  in  water, 
and  when  the  solution  is  saturated  an  equilibrium  is  established  between  gaseous 
and  dissolved  H2S.    The  solution  has  the  properties  of  a  weak  acid  which  ionizes 
in  two  stages  as  does  carbonic  acid.     Sulfides,  in  general,  are  less  soluble  than 
carbonates ;  in  the  case  of  the  alkali  metals  both  the  sulfides  and  carbonates  are 
readily  soluble,  while  in  the  case  of  the  alkali  earth  metals  the   sulfides  are 
readily  soluble  and  the  carbonates  are  not.     Questions.    Write  the  two  equations 
which  show  the  ionization  of  H2S   (dissolved  in  water)   in  two  stages.     What 
are  the  names  of  the  negative  ions  formed  ?     In  analogy  with  the  ions  of  car- 
bonic acid,  which  is  the  stronger  acid,  H2S  or  HS"?     What  are  the  principal 
substances  present  in  a  solution  of  sodium   sulfide,    Na2S?     In   a   solution  of 
sodium  hydrogen  sulfide,  NaHS?    What  reactions  take  place  when  a  strong  acid 
is  added  to  any  solution  containing  sulfide  ion?     How,  then,  will  the  solubility  of 
difficultly  soluble  sulfides  be  affected  by  the  addition  of  a  strong  acid? 

3.  Experiment.      Try  the  action  of   Na2CO3   solution   on  0.1   N  solutions   of 
Cu++,  Ag+,  and  Zn++.     In  each  case  collect  the  precipitate  on  a  filter,  wash  it 
with  water,  and  test  a  portion  for  carbonate.     Predict  the  action  of  nitric  acid 
solution  and  of  ammonium  hydroxide  on  these  precipitates,  and  test  your  pre- 
diction by  experiments. 

4.  In  the  above  experiment  with  Cu++  and  in  many  other  cases,  the  precipitate 
obtained  by  the  addition  of  Na2CO3  to  a  solution  of  a  salt  is  not  a  pure  car- 
bonate,   but    a    substance    which    contains    both    the    carbonate    and    hydroxide 
radicals.     Such  substances,  called  basic  carbonates,  are  difficult  to  obtain  pure, 
the  actual  precipitates  consisting  of  variable  mixtures  of  these  basic  carbonates, 
carbonates,   and  hydroxides. 

5.  Experiment.        Prepare    20    cc.    of    0.1    N    solutions    of  Cu++,    Ag%    and 
Zn++.    To  each  add  I  cc.  6  N  H2SO4  and  pass  in  H2S  gas*  until  the  liquid  is 
saturated.      To   determine   this,   close   the   end  of   the  test-tube   or   flask,    shake 
thoroughly  and  test  the  odor  cautiously.     Collect  the  precipitates   on   separate 
filters.     In  the  experiment  with  Zn++  and  H2S,  to  the  filtrate    (or  to  the  clear 
solution   if   there   was  no   precipitate),   add    10   cc   normal   NaAc   solution,   and 
again  saturate  with  H2S  gas.     Question.    What  effect  has  the  addition  of  Ac"  on 
the  concentration  of  H+? 

6.  H2S  and  HS~  are  weaker  acids  than  H2CO3  and  HCO3~,  respectively.    What 
can  you  conclude  about  the  relative  concentrations  of  S~~  and  CO.{~~  in  solutions 
of  H2S  and  H2CO3?    It  will  be  recalled  that 'H2CO3  does  not  precipitate  CaCO3 

*  Caution.  H2S  is  poisonous.     Work  out  of  doors  if  possible  and  do  not  breathe  the  gas. 

[481 


from  a  solution  of  CaCL  (Assignment  33,  Paragraph  6).  Question.  What 
conclusion  can  you  draw  with  regard  to  the  relative  solubility  of  CaCO3  and 
the  sulfides  prepared  in  the  preceding  Paragraph? 

7.  From  the  equation  for  the  main  reaction 

M++  +  H2S  =  MS  (solid)  +  2H+ 

it  is  evident  that  both  M++  and  H*  are  competing  for  S~~  and  that  an  increase 
in  the  H+  concentration  tends  to  reverse  the  reaction,  as  we  have  seen  in  the 
case  of  ZnS.  However  in  the  case  of  a  very  insoluble  sulfide,  CuS  for  example, 
even  though  H+  concentration  is  fairly  large  the  concentration  of  M++  will  be 
negligible  if  the  solution  is  saturated  with  H2S.  Determine  by  experiment  which 
of  these  sulfides  will  dissolve  completely  in  2  N  H2SO4.  Which  then  has  the 
greater  solubility  in  water?  Write  the  two  equations  which  show  the  mechanism 
of  the  main  reaction.  State  how  you  would  precipitate  a  sulfide  which  is  spar- 
ingly soluble  in  water,  but  many  times  more  soluble  than  ZnS  (cf.  the  precipita- 
tion of  CaCO3). 

8.  Experiment.      To     solutions     containing     Cu(NH3)4++,    Ag(NH3)o+,    and 
Zn  (  NH3 )  4++,    in    separate    experiments,    add    (NH4)2S    solution.      Repeat    the 
experiments,  using  H2S  gas  instead  of   (NH4)2S  solution.     State  in  each  case 
which  has  the  smaller  concentration  of  the  ion  of  the  metal,  a  solution  con- 
taining the   complex  ion,  or  a  solution  saturated  with  the  sulfide.     State  also 
which  is  the  less  soluble  in  water,  Cu(OH)2  or  CuS,  Ag2O  or  Ag2S,  Zn(OH)2 
or  ZnS.     From  the  fact  that  the  sulfides  are  less  soluble  than  the  carbonates, 
predict  the  action  of  H2S  on  mixtures  of  basic  copper  carbonate  and  water, 
silver  carbonate  and  water,  and  zinc  carbonate  and  water.     Test  your  answers 
by  experiment. 

9.  Experiment.     Test  the   action   of    H2S    on  solutions    of    nitric    acid    of 
different  concentrations,  0.1  N,  2.0  N,  and  6  N.  Saturate  each  solution  with 
H2S  gas  and  heat  the  mixture  almost  to  boiling.  The  white  precipitate  formed 
is  solid  sulfur  and  the  principal  reaction  is 

3H2S  +  2NO3-  +  2H+  =  3S  (solid)  +  2NO  +  4ELO. 

Questions.  What  evidence  is  furnished  by  your  experiments  that  this  is  a  slow 
reaction?  How  is  the  speed  affected  by  the  concentration  of  the  nitric  acid,  and 
by  the  temperature?  (Note.  The  products  of  the  reaction,  S,  NO,  and  H2O, 
do  not  react  with  each  other  under  any  ordinary  conditions  of  temperature  and 
pressure,  and  it  is  practically  impossible  to  realize  an  equilibrium  involving  the 
substances  shown  in  the  equation.)  Because  of  this  reaction  nitric  acid  in  a 
strongly  acid  solution  will  remove  sulfide  ion  from  the  solution  more  completely 
than  will  H+  alone  since  it  destroys  the  H2S  which  is  in  equilibrium  with  sulfide 
ion.  We  may  therefore  expect  that  nitric  acid  will  be  a  better  solvent  for 
difficulty  soluble  sulfides  than  sulfuric  or  hydrochloric  acid. 

10.  Experiment.      Collect   some   CuS  on   a   filter,   transfer   it  to   a   casserole, 
add  2  normal   HNO3   and  boil  the  mixture.       Filter,   and  test  the  filtrate   for 
Cu++.     The  residue  collected  on  the  filter  is  sulfur;  the  dark  color  is  due  to  a 
little  CuS  enclosed  in  the  sulfur.     Complete  the  equation 

3CuS  (solid)  +  2NO3-  =  3S  (solid)  +  2NO 

Problems.  (/)  Hydrogen  sulfide  is  passed  into  1  N  CuSO4  until  the  precipita- 
tion of  copper  sulfide  is  complete.  The  liquid  is  filtered  and  H2S  is  removed  by 
boiling.  What  concentration  of  H+  is  left  in  the  solution. 

(2)  Write  the  equations  for  the  neutralization  of  (a)  one  mol  H.,S  with  one  mol 
OH-,  (b)  H2S  with  excess  OH~,  (c)  HS~  with  OH~. 

(j)   Predict  the  effect    of  NH4OH    on    the    solubilities    of    CuS,    ZnS    and 

T491 


Ag2S.     Give  your  reasoning.     (If  you  try  the  experiments  use  freshly  prepared 
sulfides  moistened  with  water  which  contains  H2S.) 

(4)   Give  a  method  for  the  preparation  ot 

(a)  CuSO4  solution  from  CuS. 

(b)  Ag2S(solid)    from  AgCl. 

(c)  ZnCO3  (solid)  from  ZnS. 

(d)  Na2CO3   solution  from  Na2S   solution;  see  experiment  at  end   of 
Paragraph  8,  and  Hildebrand,  pages  207-208. 


ASSIGNMENT  38 
REVIEW  OF  THE  CHEMISTRY  OF  THE  POSITIVE  IONS  ALREADY  CONSIDERED 

1.  This  Assignment  is  inserted  in  order  that  the  student  may  summarize  and 
study  the  chemistry  of  the  positive  ions  considered   in  the  last   seven  Assign- 
ments.    If  the  student  makes  an  accurate  summary,  and  notes  carefully  similar- 
ities and  differences  in  the  results  for  the  various  ions,  he  will  find  that  it  is  not 
difficult  to  remember   the  large  number  of   facts  involved.     In  later  work,   as 
soon  as  the  chemistry  of  another  ion  has  been  investigated,  the  results  should 
be  summarized,  and  compared  with  the  results  for  the  ions  already  studied. 

2.  Summary  of  results.     Tabulate  acording  to  the  following  plan  the  results 
which  you  have  obtained  in  earlier  Assignments ;  refer  to  your  laboratory  notes 
whenever  necessary.     Write  in  a  horizontal  line  across  a  double  page  of  your 
notebook  the  names  of  the  positive  ions,   silver,   cupric   copper,   zinc,   calcium, 
sodium,  potassium  and  ammonium ;  in  the  vertical  columns  write  the  following : 

(/)   the  formulas  of  the  ions,  including  complex  ions; 
(2)   the  formulas  of  the  compounds  which  are  readily  soluble  in  water, 
(j)   the   formulas   of  the   compounds   which  are  only  moderately   soluble,  as 
Ca(OH)2,  CaS04,  Ag2SO4. 

(4)  the  formulas  of  the  insoluble  substances  (the  term  difficultly  soluble 
or  sparingly  soluble  is  preferable).  Indicate  which  is  the  least  soluble  com- 
pound of  each  metal.  When  any  ion  or  difficultly  soluble  compound  is  colored, 
give  the  color  in  the  table.  Next  make  as  general  statements  as  you  can  about 
the  solubilities  in  water  of  the  nitrates,  chlorides,  sulfates,  carbonates,  hydroxides 
(or  oxides),  and  sulfides  of  these  and  other  metals. 

3.  Show  how  you   can  apply  your  knowledge  of  the  equilibria   involved,   in 
choosing  reagents  to  dissolve  the  various  difficultly  soluble  substances  listed  in 
the  preceding  Paragraph.     In   this  connection,   for  each   complex  ion  compare 
the  concentration  of  the  ion  of  the  metal  in  a  solution  which  contains  the  complex 
ion  with  that  in  a  saturated  solution  of  the  various  difficultly  soluble  salts  of 
the   metal;    cf.    Hildebrand's    discussion   of   the   reactions   of    silver   ion,   pages 
204-205,  and  the  following  example.     Since  ZnCO3  dissolves  readily  in  excess 
NH4OH   solution,  and  ZnS  does  not,  it  may  be  concluded  that  in  a  solution 
which  contains  Zn  (NH3)4++  and  NH4OH  the  concentration  of  Zn++  is  less  than 
in  a  saturated  solution  of  ZnCO3  but  greater  than  in  a  saturated  solution  of  ZnS. 

4.  The  Summary  in  Paragraph  2  was  obtained  from  the  results  of  experiments 
on   the   reactions   between   the   positive   ions   and   various   reagents.     Conversely 
the  results  of  such  experiments  can  be  predicted  by  means  of  the  Summary  and 
a  knowledge  of  the  various  equilibria  involved. 

[50] 


5.  Table  of  Reactions.     Prepare  for  future  reference  a  table  of  the  reactions 
between  dilute  solutions  of  the  positive  ions  considered  above  and  the  following 
reagents : 

(a)  HC1  or  a  soluble  chloride, 

(b)  NaOH  in  small  amount, 

(c)  NaOH  in  excess, 

(d)  NH4OH  in  small  amount, 

(e)  NH4OH  in  excess, 
(/)   H2C03  or  CO,  gas, 
(g]   a  soluble  carbonate, 

(A)  H,S  in  0.3  N  H+  solution, 

(*)  a  soluble  sulfide  as  Na2S,  or  (NH4)2S  in  NH.tOH  solution, 

(/  )  sodium  cobaltinitrite  in  the  presence  of  a  few  drops  of  HAc. 

When  no  reaction  takes  place  mark  X  at  the  proper  place  in  the  table.  In 
other  cases  write  the  formulas  of  the  substances  formed,  note  the  colors,  and 
mark  the  precipitates  by  underlining  the  formulas.  Be  sure  that  you  can 
write  the  equation  for  each  reaction.  Try  experiments  whenever  it  is  neces- 
sary, e.  g.  when  your  laboratory  notes  are  incomplete  or  seem  to  be  incorrect, 
and  when  there  is  a  blank  space  in  your  table. 

6.  Test  your  knowledge  of  the  reactions  just  considered  by  writing  out  from 
memory  portions  of  the  Table  of  Reactions,  Paragraph  5  ;  write  equations  for  the 
reactions  that  take  place,  and  discuss  briefly  the  equilibria  involved. 


Note.  It  is  recommended  that  the  student  proceed  with  Assignment  51,  the 
first  in  the  Section  entitled  Systematic  Qualitative  Analysis.  It  is  important  that 
the  student  become  thoroughly  familiar  with  the  reactions  already  considered 
before  proceeding  to  study  other  reactions,  and  this  Assignment  deals  only  with 
the  ions  already  considered. 


51 


SECTION  IV 
REACTIONS  OF  IONS,  CONTINUED 


ASSIGNMENT  41 
OXIDATION  AND  REDUCTION.  REPLACEMENT  REACTIONS.  ELECTRICAL  CELLS 

References.     Hildebrand,  Chapter 'VI  pages  93-94,  Chapter  V  page  70,  Chapter 

XV  pages  234-237. 

1.  In  Assignment  41  and  the  three  following  Assignments  we  shall  consider 
examples  of  oxidation  and  reduction  reactions.     In  these  reactions  changes  of 
valence  occur,  and  they  are  therefore  easily  distinguished  from  other  reactions 
such  as  the  formation  of  weak  electrolytes  or  the  precipitation  of  difficultly  soluble 
salts.     The  following  reactions  can  now  be  recognized  as  oxidation  and  reduc- 
tion reactions : 

Zn  (solid)  +  2H+  =  Zn++  +  Ho  (Assignment  3) 

H,  +  CuO  (solid)  =  Cu  (solid)  +  HoO  (Assignment  4) 

Ca  (solid)  +  2H2O  =  Ca+V  +  2OH-  +  H,  (Assignment  32) 

3H2S  +  2NO3-  +  2H+  ==  3S  +  2NO  +  4EUO  (Assignment  37) 

3d-  +  NO3-  +  4H+  =  3/2  C12  +  NO  +  2H.,O 

=  C12  +  NOC1  +  2H26  (Assignment  34) 

Point  out  the  changes  of  valence  in  each  reaction. 

2.  In  the  first  of  the  above  reactions  the  oxidation  of  Zn  to  Zn++  and  the 
reduction  of  2H+  to  H2  consist  simply  in  the  transfer  of  two  electrical  charges 
from  two  hydrogen  atoms  to  one  zinc  atom   (which  in  reality  is  the  transfer 
of  two  electrons  from  the  zinc  atom  to  the  two  hydrogen  atoms).      In  the  pres- 
ent Assignment  we  shall  study  a  number  of  examples  which  are  characterized 
by  such  an  interchange  of  electrical  charges.  The  reaction  between  copper  and 
silver  ion  will  first  be  studied  quantitatively  and  then  a  number  of  other  cases 
will  be  taken  up  qualitatively. 

3.  Experiment.     Obtain  at  the    office  100  cc.  0.10  N  silver    nitrate,    and    a 
special  filter  paper  which  burns  to  give  an  ash  of  negligible  weight.     Clean  a 
straight  piece  of  thin  copper  wire  and  weight  about  0.3  g.  to  5  mg.     Cover  the 
copper  wire  with  100  cc.  of  0.10  N  silver  nitrate  solution  and  let  it  stand  until 
the  copper  has  disappeared.    The  precipitate  formed  is  metallic  silver.     Questions. 
What   substance   is   responsible   for  the   color  of  the  final   solution?     What  is 
the  color  of  silver  ion?     Heat  a  porcelain  crucible,  let  it  cool,  and.  weight  to  5 
mg.     Collect  the  silver  by  filtering  through  the  special  filter  paper  and   wash 
the  precipitate  until  the  washings  are  colorless.     (Save  the  filtrate,  and  place  in 
it  another  piece  of  copper  in  order  to  recover  the  rest  of  the  silver.)     Fold  the 
filter  paper  so  that  it  will  fit  into  the  crucible,  heat  with  a  small  flame  until  it  is 
dry,  and  then  with  a  larger  flame  until  the  paper  is  completely  consumed.     Let 
the  crucible  cool  and  again  weigh  to  5  mg.    Place  all  the  silver  in  the  jar  marked 
"Silver  Waste." 

4.  Calculations.     Calculate  (a)    the   weight   of    silver    formed    in   the   experi- 
ment; (b)  the  weight  of  silver  precipitated  by  one  gram  atom  of  copper.     Com- 
pare the  latter  number  with  the  atomic  weight  of  silver.      (The   formula  of 
silver  ion  is  Ag+.)     What  is  the  formula  of  copper  ion? 

5.  Questions.     Give  the  formulas    of    silver    nitrate,   silver    sulfate,    copper 
nitrate  and  copper  sulfate.     If  one  gram  molecule  of  silver  ion  has  associated 
with  it  96,500   coulombs  of  positive  electricity  what   quantity   of   electricity   is 

[521 


associated  with  one  gram  molecule  of  copper  ion?  With  a  gram  molecule  of 
nitrate  ion?  Write  the  equation  for  the  action  of  silver  ion  on  copper.  How 
would  the  result  of  this  experiment  have  been  influenced  by  replacing  the  silver 
nitrate  with  silver  sulfate  solution? 

6.  Experiment.    Into  10  cc.  N  Cu(NO3)2  place  a  piece  of  lead. 

7.  Repeat  the  experiment    with  a    solution  of    lead    nitrate,    Pb(NO3)2,  and 
zinc. 

8.  Write  the  equations  for  all  the  reactions  between  metals  and  positive  ions 
which   you  have   studied    (Assignments   3  and   41).      Questions.    The  negative 
ion  present  in  each   experiment  is  not  shown   in   the   equation;   what  experi- 
mental fact  is  thereby  implied?     If  in  one  of  the  above  experiments  the  dry 
salt  had  been  used  instead  of  its  solution,  would  the  reaction  occur  sufficiently 
rapidly  at  room  temperature  to  be  detected?     Try  an  exepriment,  if  necessary, 
with  lead  and  solid  copper  sulfate. 

9.  Experiment.       Determine    by    experiment    which    of    the    metals    used    in 
this  Assignment  will  dissolve  readily  in  dilute  hydrochloric  acid    (or  sulfuric 
acid).* 

10.  Arrange    the    metals    used    in    these    experiments    in  the  order  in  which 
they  will  replace  each  other  from  solution.     After  referring  to  your  lecture 
notes  state  where  hydrogen  belongs  in    this    replacement    series.     From    this 
replacement  series  predict  what  will  happen  when, 

Silver  nitrate  solution  is  treated  with  lead. 
Copper  nitrate  zinc. 

Copper  nitrate  silver. 

Experiment.  Test  the  correctness  of  your  prediction. 

11.  Replacement  series  for  the  halogens.      Experiment.       Make   a   small   gas 
generator   (see  Assignment  33,   Paragraph  3)   and   charge  the  flask  with  about 
2    g.    of    manganese    dioxide.       Pour    5    cc.   concentrated    HC1  into  the  flask 
and  warm  gently  until  a  continuous  stream  of  chlorine  is  evolved.     Pass  a  few 
bubbles  of  the  gas  through  a  very  dilute  solution  of  potassium  bromide,  pre- 
pared  by  diluting  a   few   drops   of  the   laboratory   solution  with    5    cc.    water. 
Shake  the  solution  with  I  cc.  of  carbon  bisulfide.     The  brown  color  in  the  car- 
bon bisulfide  layer  is  characteristic  of  bromine. 

12.  Experiment.     To  1  cc.  potassium  iodide  solution  add  a  drop  of  bromine 
water.     Shake  with  carbon  bisulfide.    The  violet  color  is  characteristic  of  iodine. 

13.  Questions.     In  each  of  the  above  experiments  write  the  equation  for  the 
reaction  which  has  taken  place.    Point  out  the  changes  of  valence  and  the  inter- 
change of  electrical  charges.  Predict  what  reaction  will  occur  when  (a)  chlorine 
gas  is  passed  through  KI  solution,   (b)  bromine  is  added  to  KC1  solution,  and 
(c)   iodine  is  added    to  KBr    solution.        Experiment.       Test    the    correctness 
of  your  predictions. 

14.  The  Conversion  of  Chemical  into  Electrical  Energy.    The  reactions  studied 
in  the  present  Assignment  may  be  utilized  to  convert  chemical  into  electrical 
energy.     In  each  of  the  reactions  an  interchange  of  electrical  charges  has  taken 
place;  and  the  problem  in  the  construction  of  an  electrical  cell  is  to  arrange  the 
substances   involved  in   the  reaction   in   such   a   way   that   this   interchange   of 
electricity  can  take  place  only  through  an  external  conductor  such  as  a  wire 
joining  the  electrodes.     Thus,  in  the  reaction  between  zinc  and  copper  ion  two 
changes  of  an  electrical  nature  are  involved:  on  the  one  hand  zinc  is  changed 
to  zinc  ion,  and,  on  the  other,  copper  ion  is  changed  to  copper.    When  the  zinc 

*The  action  of  nitric  acid  on  metals  will  be  studied  in  Assignment  42. 

[531 


is  placed  directly  in  a  solution  of  a  copper  salt  the  electrical  changes  both  occur 
at  the  surface  of  the  zinc  and  the  transfer  of  electricity  cannot  be  directly 
detected;  but  when  we  arrange  the  materials  so  that  the  zinc  electrode  dips  into 
a  zinc  sulfate  solution  contained  in  a  porous  cup,  and  this  is  placed  in  a  vessel 
containing  copper  sulfate  solution  and  the  copper  electrode,  then  no  reaction 
takes  place.  If  now  we  join  the  two  electrodes  by  a  wire,  electricity  passes 
through  the  wire  and  zinc  will  change  to  zinc  ion  at  the  surface  of  the  zinc, 
while  copper  ion  changes  to  copper  at  the  surface  of  the  copper.  The  total 
change  in  the  cell  is  of  course  the  same  as  when  zinc  acts  directly  on  copper 
ion. 

15.  Draw  a  diagram  of  the  elecrical  cell  described  above.     Question.       What 
substances   carry   the   electricity   from   one   electrode   to   the   other   through   the 
solution  ? 

16.  In   a   similar   way   the   reaction   between   chlorine   and   bromide   ion   may 
be  used  to  generate  electrical  energy;  but,  to  lead  the  current  into  and  out  of 
the  solution,  we  must  use  some   substance,  as  graphite,   which  conducts  elec- 
tricity and  is  not  attacked  by  chlorine  or  bromine.     In  one  part  of  this  elec- 
trical  cell  there  would  thus  be  a  graphite  rod  dipping  into  a  solution  which 
contains  bromide  ion  and  bromine,  and  in  the  other  part  a  graphite  rod  in  a 
solution  of  chloride  ion  and  chlorine.     The  two  solutions  must  be  separated, 
e.  g.,  by  a  porous  cup,  in  order  to  prevent  the  chlorine  from  reacting  directly 
with  the  bromide  ion.     When  the  graphite  electrodes  are  joined  by  a  wire  a 
current  will  flow,  and  at  the  one  electrode  bromide  ion  will  change  to  bromine 
while  at  the  other  electrode  chlorine  will  change  to  chloride  ion. 

17.  Problems.      (/)      What    quantity    of    electricity     (in     coulombs)     flows 
through  the  circuit  of  the  cell  described  in  Paragraph  13  when  one  mol  of  zinc 
disappears?     What  weight  of  copper  will  at  the  same  time  be  deposited  on  the 
copper  electrode? 

(2)  Sketch  the  arrangement  of  a  cell  in  which  the  reaction  between 
metallic  copper  and  silver  ion  is  used  to  produce  an  electric  current.  When  a 
current  is  taken  from  this  cell  what  changes  take  place  at  each  electrode? 

(j)  How  many  coulombs  of  electricity  could  be  obtained  by  the  transforma- 
tion of  1  liter  chlorine  gas  (measured  at  0°  C  and  760  mm.)  to  chloride  ion? 


ASSIGNMENT  42 

OXIDATION  OF  METALS  TO  THEIR  IONS.  TABLE  OF  OXIDIZING 
AND  REDUCING  AGENTS 

Reference.    Hildebrand,  Chapter  XV  pages  226-245. 

1.  In    Assignment    41    we    found    that    certain    metals    can    be    oxidized    by 
hydrogen  ion  and  that  other  metals  cannot.    We  may  say,  therefore,  that  certain 
metals  are  sufficiently  strong  reducing  agents  to  reduce  hydrogen  ion  to  hydro- 
gen and  that  others  are  not,  and  it  follows  that  any  reagent  which  will  change 
one  of  the  latter  metals  to  its  ion  is  a  more  powerful  oxidizing  agent  than  is 
hydrogen  ion.     In  the  present  Assignment  we  shall  study  the  action  on  zinc, 
copper,  and  silver  of  two  such  oxidizing  agents,  nitric  acid  and  bromine. 

2.  Arrange    copper,    silver,    zinc    and    hydrogen,    in    a   vertical    column,    in    a 
replacement  series.     Beside  each  symbol  write  the  formula  of  the  ion  to  which 
the    element    can    be    oxidized.      Your    table    now    is    a    table    of    oxidizing 
and   reducing   agents   arranged   according   to   their   strengths,    and    this    table 

[541 


may  be  extended  to  include  other  substances  than  the  metals  and  their 
ions.  Questions.  Which  is  the  more  powerful  reducing  agent,  Zn  or  Cu?  Cu 
or  Ag?  Which  is  the  more  powerful  oxidizing  agent,  Ag+  or  Cu++?  Ag+  or  H+? 
(State  the  experimental  facts  upon  which  you  based  your  answers  to  the  pre- 
ceding questions.)  In  describing  the  reaction  between  zinc  and  dilute  hydro- 
chloric or  sulfuric  acid  solution  is  it  correct  to  say  that  metallic  zinc  has  com- 
bined with  chloride  ion  or  with  sulfate  ion?  How  can  the  solid  salts  be  pre- 
pared from  the  solutions  obtained  in  such  experiments? 

3.  Action  of  nitric  acid  on  zinc,  copper,  and  silver.        Experiment.         Treat 
about   l/±   grams  of  zinc  with  2  .V  HNO3   in  a   small  beaker,  covered  with  a 
watch-glass.     If  the  reaction  is  slow  warm  the  mixture.     The  main  reaction 
under  these  experimental  conditions  is 

3Zn  (solid)  +  2NO3-  +  8H+  =  3Zn+*  +  2NO  (gas)  +  4H2O. 

4.  Repeat  the   experiment  with    about    0.5    g.    copper,    and    with    a    small 
piece    of    silver.      The    reactions    are    similar    to    the    one    between    zinc    and 
nitric  acid,  and  the  change  from  Zn  to  Zn++  corresponds  exactly  to  the  change 
from  Cu  to  Cu++  and  from  2  Ag  to  2  Ag+.    Write  the  equations  for  the  reactions. 
Question.     How  is  the  speed  of  the  reaction  between  nitric  acid  and  the  metals 
affected  by  a  change  in  temperature  and  by  a  change  in  the  concentration  of  the 
acid? 

5.  Action   of    bromine    on   zinc,    copper  and   silver.       The  halogens,  chlorine, 
bromine  and   iodine,   all   have  a  similar  action  on  metals.     To   illustrate  this 
action  we  use  bromine  because  its  aqueous   solution  is  a   convenient  reagent. 
The  solution  of  chlorine  in  water  does  not  keep  well,  while  iodine  is  a  solid 
which  is  only  very  slightly  soluble  in  water.     Experiment.     Treat  a  small  piece 
of  zinc  with  bromine  water.     After  a  few  minutes  pour  the  solution   into   a 
porcelain  dish,  and  heat  it  to  boiling  to  expel  any  bromine  that  is  present.    To 
a  portion  of  the  solution  add  a  few  drops  of  silver  nitrate:    AgBr,  like  AgCl 
and  Agl,  is  a  difficulty  soluble  salt.     Test  another  portion  of  the  solution  for 
zinc  ion.     What  evidence  have  you  that  a  reaction  takes  place  between  zinc 
and  bromine?     Repeat  the  experiment  with  copper  and  bromine  water.    Silver, 
also,   reacts  with  bromine. 

6.  The  equation  for  the  reaction  between  zinc  and  bromine  solution  is 

Zn  (solid)  +  Br2  =  Zn++  +2Br 

Since  in  this  reaction  zinc  is  oxidized  to  zinc  ion  we  must  conclude  that  bromine 
is  reduced  when  it  changes  to  bromide  ion.  Write  the  equation  for  the  action 
of  bromine  on  copper  and  on  silver. 

7.  Review    the    replacement  series  for  the    halogens     (Assignment    41)     and 
arrange  chloride  ion — chlorine,  bromide  ion — bromine,  iodide  ion — iodine  in  a 
table  so  as  to  give  the  relative  strengths  of  the  halogens  as  oxidizing  agents  and 
of  the  halide  ions  as  reducing  agents.    Add  this  table,  in  its  proper  place,  to  the 
one  previously  prepared  for  the  metals  and  their  ions.     Place  Ca  —  Ca+*  and 
Na — Na+  in  this  table.     Show  your  completed  table  at  once  to  the  instructor. 

8.  We  can  determine  the  position  of  nitric  acid  in  this  table  of  oxidizing  and 
reducing  agents  by  reference  to  the  results  of  the  experiment  in  Paragraph  6, 
Assignment  34,  when  hot  concentrated  nitric  acid  was  used  to  remove  chloride  ion 
from  a  solution.    Chlorine  was  formed  and  we  can  therefore  decide  that  hot  con- 
centrated nitric  acid  is  a  stronger  oxidizing  agent  than  chlorine.   Question.   Is  con- 
centrated nitric  acid  a  stronger  oxidizing  agent  than  bromine  or  iodine? 

9.  Study  the  method  of  balancing  oxidizing  and  reducing  reactions.     Hilde- 
brand,  pages  226-233. 

10.  Problems,     (i)     Summarize  the  behavior  of  Na,  K,  Ca,  Zn,  Cu,.  and  other 

[55] 


metals  discussed  in  the  lectures  by  dividing  the  metals  into  the  three  following 
classes : 

I.     Those  which  react  readily  with  water.     Hydrogen  is  evolved. 
II.     Those  which   react   readily  with   HC1  or  H2SO4   solution,    but    not 

with  water.     Hydrogen  is  evolved. 

III.     Those  which  dissolve  readily  in  HNO3  or  Br,  solution,  but  not  in 
HC1,  H2SO4  or  water.     H2  is  not  evolved. 

Write  the  equation  for  the  action  of  HC1  or  H2SO4  solution  on  metallic  sodium. 

(2)  In  the  experiment  in  Paragraph  11,  Assignment  41,  manganese  dioxide, 
MnO2,  reacted  with  chloride  ion  in  acid  solution  to  form  chlorine.  Manganous 
ion,  Mn++,  was  formed.  Write  the  equation  for  the  reaction.  Predict  what  will 
happen  when  copper  is  treated  with  MnO2  in  the  presence  of  dilute  sulfuric 
acid. 

(j)  Sketch  the  arrangement  of  a  cell  in  which  the  reaction  between  zinc 
and  bromine  is  used  to  produce  an  electric  current.  What  changes  take  place 
at  the  electrodes?  State  how  you  could  bring  about  the  reversal  of  all  the  reac- 
tions in  this  cell. 


ASSIGNMENT  43 
FERROUS  AND  FERRIC  IONS 


Reference,  in  this  and  later  Assignments :      A  standard  text  on  Inorganic  Chem- 
istry.   Note.    Read  again  Asignment  31,  Paragraph  2,  and  Assignment  38. 

1.  In  Assignment  43   we   shall   study   the   chemistry  of   iron  as   an   example 
of    a    metal    which    forms    two    series    of    compounds   in  which    the    valence 
of  the  metal  is  different.     You  will  find  on  the  laboratory  shelves  solutions  and 
solid  salts  of  ferrous  and  of  ferric  iron.     Give  the  valences  of  the  following: 
sulfate  ion;  sulfate  in  ferrous  sulfate,  FeSO4;  Fe   (metal);  ferrous  ion,  Fe++; 
ferric  ion,  Fe+++;  iron  in  a  ferrous  salt,  and  in  a  ferric  salt.    Write  the  formulas 
of  the  chlorides,  nitrates,  hydroxides,  and  oxides  of  ferrous  and  of  ferric  iron. 

2.  Many  double  salts  are  found  among  iron  compounds. 
FeSO4'(NH4)2SO4' 6H2O,  ferrous  ammonium  sulfate,  and 

Fe2(SO4)3 '(NH4)2SO4' 24H2O,  ferric  ammonium  sulfate  (ammonium 
iron  alum)  are  examples  of  double  salts  whose  solutions  dissociate  into  the  ions 
of  the  simple  salts.  These  two  salts  may  be  used  in  studying  the  properties  of 
Fe++  and  Fe+++,  respectively.  Give  the  formulas  of :  alum,  potassium  iron 
alum,  sodium  iron  alum. 

3.  The     ferrocyanides     and     ferricyanides     are     examples     of     double     salts 
whose  solutions   contain   complex  ions   at   high   concentrations.     The   formula, 
Fe(CN)2  •  4NaCN,   for   sodium   ferrocyanide   shows   its   relation   to   the   simple 
salts.     It  can  be  prepared  by  dissolving  the  difficultly  soluble   ferrous  cyanide, 
Fe(CN)2,  in  excess  of  NaC'N,  but  in  a  solution  of  the  pure  salt  the  concentra- 
tions of  Fe++  and  CN~  are  extremely  small.     Accordingly  the  formula  is  usually 
written  Na4Fe(CN)6  to  show  that   the  ferrocyanide   radical  Fe(CN)0  is   the 
negative    constituent    of    the    salt.      Similarly    Na3Fe(CN)6    is    the    formula 
usually   written    for   sodium    ferricyanide.      Questions.      What    is    the   valence 
of    the    ferrocyanide    radical?      Of    the    ferricyanide    radical?      What    is    the 
formula  of  potassium  silver   cyanide?     How   is  it  prepared?     What  ions   are 
present  in  its  solution?     Caution.    Soluble  cyanides  and  hydrocyanic  acid  are 
deadly  poisons :   do  not  try  unnecessary  experiments   with   them.     Ferro   and 
ferricyanides  are  not  dangerous. 

[561 


4.  Color  tests    for  Fe++  and    Fe+++.     Experiment.     To    a    dilute    solution    of 
FeCl3  add  a  drop  or  two  of  K4Fe(CN)6  solution.     Repeat  with  K3Fe(CN)8. 
Prepare  a  solution  of  a  ferrous  salt  by  dissolving  a  small  amount  of  FeSO4 
or  of  FeSO4*(NH4)2SO4  in  a  little  dilute  sulfuric  acid,  and  test  portions  of  this 
solution    at    once    with   K4Fe(CN)c   and    with    K3Fe(CN)6.       The     character- 
istic  dark   blue   precipitates,    or   colloidal   solutions,   obtained   in   two   of   these 
four  experiments  are  probably  identical.     Equations  may  be  omitted  for  these 
reactions. 

5.  Oxidation    of   Fe++  to  Fe+++.       Experiment.       Test  for  Fe+++  an  acid  solu- 
tion of  FeSO4   which   has   been   standing  In  the  laboratory   for   several   days. 
The  oxygen  of  the  air  slowly  oxidizes  Fe++  to  Fe+++ 

4Fe++  +  O2  +  4H+  =  4Fe+++  +  2H2O. 

Experiment.  To  a  few  drops  of  a  ferrous  salt  solution  in  a  porcelain  dish  add 
a  little  HNO3,  heat  the  mixture,  dilute  with  water,  and  test  whether  the  solu- 
tion now  contains  Fe"1"1"  or  Fe+++.  Try  an  experiment  with  another  oxidizing 
agent  and  determine  whether  or  not  it  oxidizes  Fe++  to  Fe+++. 

6.  Note.    Write  equations  for  all  reactions. 

7.  Reaction   between  Fe  and  Fe+++.      Experiment.      Place    some    iron    wire 
or  iron  filings  in  a  very  dilute  solution  of  FeCl3,  stir  the  mixture,  and  test 
small  portions  of  the  solution  from  time  to  time  for  Fe++  and  Fe+++. 

8.  Reactions  of  iron  ivith  H+,  and  ivith  Cu^.     Experiment.     Try  the  action 
of  dilute  HC1  or  H2SO4  on  a  small  piece  of  iron;  heat  the  mixture,  if  neces- 
sary.    Predict  whether  ferrous  or  ferric  ion  is  formed,  and  test  the  correctness 
of  your  answer.     Experiment.    Cover  a  small  piece  of  iron  with  a  solution  of 
a   copper   salt. 

9.  From  the  results  of  the  above  experiments,  or  from  any  other  source  of 
information,  determine  approximately  the  position  of  Fe — Fe++  and  of  Fe++ — Fe+++ 
in  your  table  of  oxidizing  and  reducing    agents,    Assignment    42,    Paragraphs 
2  and  7.     It  is  not  advisable  to  attempt  to  assign  a  definite  place  in  one  large 
table  to  each  of  the  oxidizing  and  reducing  agents  which  you  will  study.     Thus 
in  the  case  of  Fe++— Fe+++  it  is  sufficient  to  note  that  Fe++— Fe+++,  I~— 12  and 
Ag — Ag+  are  in  the  same  part  of  the  table:  Fe++,  I~  and  metallic  silver  are  all 
readily  oxidized      by  powerful  oxidizing  agents,  while  Fe+++,   I2  and  Ag+  are 
easily  reduced  by  powerful  reducing  agents.    Questions.    Which  is  the  stronger 
reducing   agent,    Fe"1"1"   or   Fe?    Which   is    the   stronger   oxidizing   agent,    Fe*++ 
or  Fe++? 

10.  Reduction  of  Fe+++  by  H2S  in  acid  solution.       Experiment.       To    a    few 
drops  of  FeCL  solution  add  1  cc.  6  N  HC1  and  about  20  cc.  water,  and  saturate 
the  solution  with  H2S  gas.     Boil  the  mixture  in  a  porcelain  dish  to  coagulate 
the  precipitated  sulfur  and  expel  the  H2S,  filter  and  at  once  test  a  portion  of  the 
clear  solution  for  Fe+++.   Treat  the  remainder  of  the  solution  again  with  H2S  if 
Fe+++   seems   to   be   present.      Name   another   reducing   agent   which   might   be 
expected  to  reduce  Fe+++  to  Fe++  and  try  the  experiment. 

11.  Other  reactions  of  Fe+++.    Experiments.     Treat  small  quantities  of  dilute 
solutions  of  FeCl3  with  the  following  reagents: 

(a)  NH4OH  in  small  amount  and  in  excess.    The  precipitate  is  Fe(OH)s. 

(b)  NaOH  in  small  amount  and  in  excess. 

(c)  Na2CO3   or    (NH4),CO3.   Filter;   acidify  the  filtrate  and  test  it  for  Fe+++ 
with   K4Fe(CN)6;   wash  with   water  the  precipitate    (obtained   by   adding 
carbonate),  and  test  whether  it  is  a  carbonate  or  hydroxide. 

(rf)   NH4OH  or  NaOH  and  then  a  sulfide  as  H2S  or   (NH4)2S.     The  black 
precipitate  is  mainly  Fe2S3,  but  may  contain  some  FeS  and  S. 

[57] 


12.  Other  reactions  of  Fe++.     Repeat    the    experiments    in    (11)    with    small 
quantities   of   dilute   solutions   of   ferrous   salt.        In    (a)    and    (b)    note   what 
happens   when   moist   ferrous   hydroxide   is   exposed   to    the   air;   the    oxygen 
reacts  much  more  rapidly  with  moist  Fe(OH)2  than  with  Fe++  in  acid  solution. 
The  brown  compound  formed  in  the  test  for  nitrate  ion,  Assignment  31,  Par- 
agraph 13,  is  an  addition  compound  of  NO  and  ferrous  ion,  probably  Fe(NO)++. 
This  compound  is  formed  when  NO  gas  is  passed  into  a  cold  ferrous  salt  solution, 
and  when  nitric  acid,  nitrous  acid,  or  NO2  is  reduced  by  excess  Fe++.     Questions. 
What  happens  when  a  solution  which  contains  this  addition  compound  is  heated? 
Is  the  rate  of  reduction  of  nitrate  ion  in  dilute  acid  solution  by  Fe++  a  rapid  or  a 
slow  reaction?     Experiment.     Determine  whether  the  reduction  of  nitrous  acid, 
HNO2,  is  a  rapid  or  a  slow  reaction  by  adding  to  10  cc.  water,  2  cc.  FeSO4 
solution,  a  few  drops  oi  6  N  H2SO4,  and  a  drop  of  a  nitrite  solution. 

13.  Summarize   in   a   table    for   future   reference   the    reactions   of    Fe++   and 
Fe+++  with  the  following  reagents:  a  strong  oxidizing  agent,  as  hot  moderately 
concentrated  HNO3;  a  strong  reducing  agent,  as  zinc  in  acid  solution;  H2S 
in  0.3  N  acid  solution,  and  the  reagents  used  in  Paragraphs  11  and  12.    Below  the 
table  give  methods  of  dissolving  the  difficultly  soluble  salts  of  iron  which  you 
have  studied. 

14.  Problems.     Try  experiments  when  necessary. 

(/)   How  would  you  transform 

(a)  ferric  chloride  to  ferric  sulfate; 

(b)  ferric  sulfate  to  ferric  chloride; 

(c)  ferrous  sulfate  to  ferric  sulfate; 

(d)  ferric  sulfate  to  ferrous  sulfate? 

(2)  Fe(OH)2  is  moderately  soluble  in  a  solution  of  an  ammonium  salt, 
as  NH4C1,  and  Fe(OH)3  is  not.  Outline  an  experiment  to  demonstrate  this. 
Which  is  the  less  soluble  in  water,  Fe(OH)2  or  Fe(OH)3? 

(j)  Suggest  a  method  of  preparing  Fe(OH)3  and  ZnS  from  a  single  portion 
of  a  solution  which  contains  Fe++,  Zn++,  H+  and  SO4~~. 


ASSIGNMENT  44 
MERCUROUS  AND  MERCURIC  IONS 


1.  In  the  present  Assignment  we  shall  study  the  chemistry  of  another  metal 
which  forms  two  series  of  compounds.    You  will  find  on  the  laboratory  shelves 
solutions  of  mercurous  nitrate,  HgNO8;  and  of  mercuric  nitrate,  Hg(NO3)2. 
What  is  the  valence  of  mercury  in  a  mercurous  salt;  in  a  mercuric  salt?   Write 
the  formulas  of  the  chlorides,  sulfates  and  oxides  of  mercurous  and  of  mer- 
curic mercury. 

2.  Mercurous    chloride    and    mercuric    chloride.       Experiment.       To    1     cc. 
N  HgNO3  add  20  cc.  water  and  1  cc.  6  N  HC1.     Repeat  the  experiment  with 
N  Hg(NO3)2,  and  give  a  method  of  distinguishing  between  soluble  mercurous 
and  mercuric  salts.     Note.     Mercuric  chloride  is  one  of  the  few  examples  of  a 
weak  salt.    The  concentrations  of  Hg++  and  of  Cl~  in  its  solution  are  very  small. 

3.  Reduction  of  mercuric  ion  in  stages.     Tin  forms  two  series  of  compounds 
stannous  and  stannic,  in  which  the  valence  of  tin  is  -\-  2  and  +  4  respectively. 
Stannous  ion,  Sn++,  is  a  good  reducing  agent  since  it  is  easily  oxidized  to  Sn+++*. 
The  chlorides  of  tin  are  both  soluble  in  water;  and  the  common  stannous  salt 
laboratory  reagent  is  SnCl2  solution  to  which  HC1  has  been  added  to  prevent 

[581 


the  precipitation  of  a  basic  salt.  Experiment.  To  1  cc.  N  Hg(NO3)2  and 
20  cc.  water  add  one  drop  SnCL  solution.  Divide  the  mixture  into  two  portions 
and  to  one  portion  add  SnCl2  until  it  is  present  in  excess.  Observe  the  differ- 
ence in  appearance  of  finely  divided  mercury  and  a  large  drop  of  the  metal. 
Suggest  a  method  of  testing  a  solution  for  the  presence  of  mercuric  mercury. 
Try  your  method  with  a  dilute  solution  of  HgCl2. 

4.  Reaction    between    mercuric    ion    and    mercury.     Experiment.     To   1    cc. 
N  Hg(NO3)2  add  a  drop  of  mercury,   shake  the  mixture  gently,  and  after 
several  minutes  add  20  cc.  water.     Pour  off  the  solution  and  test  it  for  Hg+ 
with  HC1.     Filter  off  the  mercurous  chloride  and  test  the  filtrate  for  mercuric 
mercury.     If  there  is  more  than  a  mere  trace  repeat  the  experiment  and  leave 
the  mercury  in  contact  with  the  solution  for  a  long  time.     The  reaction  be- 
tween Hg  and  Hg++  is  similar  to  that  between  Fe  and  Fe+++.     In  each  case,  in 
the  presence  of  excess  of  the  metal,  the  reaction  will  proceed  until  at  equi- 
librium the  concentration  of  the  ion  of  higher  valence,  Hg++  or  Fe+++,  is  ex- 
tremely  small.      Questions.    Which   is   the  stronger  oxidizing  agent,   Kg**   or 
Hg+?    Which  is  the  stronger  reducing  agent,  Hg+  or  Hg? 

5.  Reduction   of  Hg+  and  oxidation  of  Hg.     By  experiments  with  some  of 
the  metals  you  have  studied  and  small  amounts  of  HgNO3  solution  determine 
approximately  the  position  of  Hg  —  Hg*  in  your  table  of  oxidizing  and  reduc- 
ing agents.     In  these  experiments  any  metallic  mercury  formed  can  usually  be 
seen  on  the  surface  of  the  metal,  but  the  solution  may  also  be  tested  for  the 
dissolved  metal  after  precipitating  the  mercurous  mercury  with  HC1.     From 
your  result  predict  whether  metallic  mercury  will  dissolve  in  dilute  sulfuric 
acid  ,  and  test  your  prediction.    Experiment.     Treat  a  very  small  drop  of  mer- 
cury with  about  30  cc.  2  N  HNO3  in  a  porcelain  dish  or  beaker.     If  the  reac- 
tion is  slow  heat  the  mixture.     While  the  reaction  is  in  progress  pour  a  small 
amount  of  the  solution  into   10  cc.   water  and  test  for  Hg+  and  Hg++.     When 
a  moderate  amount  of  the  metal  has  dissolved,  pour  off  the  solution  for  use 
in   the   next   experiment. 

6.  Oxidation  of  Hg+  and  of  HgCl.       Experiment.       Continue  the  preceding 
experiment  to  determine  how   rapidly  hot  2  N   HNCX  oxidizes  Hg+.     Repeat 
the  experiment  with  hot  concentrated  HNO3  and  a  small  quantity  of  HgNO3. 
Experiment.     Collect  some  HgCl  on  a  filter.    Treat  a  small  portion  in  a  porce- 
lain dish  with  hot  2  N  HNO3,  and  another  portion  with  hot  2  N  HNO3,  to 
which  some  HC1  has  been  added.     Repeat  the  experiments,  if  necessary,  with 
more  concentrated  HNO3,  and  note  whether  the  speed  of  oxidation  is  increased 
by  the  presence  of  HC1.     Experiment.     Treat  a  small  portion  of  HgCl  with 
bromine  water,  and  test  the  solution  as  before  for  mercuric  mercury. 

7.  Difficulty    soluble    salts    of    mercuric    mercury.       Experiment.       Test  the 
action  of  each  of  the  following  reagents  on  a  dilute  solution  of  Hg(NO3)2.  Save 
the  precipitate  for  use  in  (8). 

(a)  H2SO4  or  a  soluble  sulfate. 

(b)  Na^COg.     The  precipitate  is  a  basic  carbonate. 

(c)  NaOH.     The  precipitate  is  HgO. 

(d)  NH4OH.     When    a    chloride    is    present    the  precipitate  is    Hg(NH2)Cl, 
ammonobasic   mercuric   chloride,   which   is   closely   related   to   Hg(OH)Cl, 
basic  mercuric  chloride.     Analogous  but  more  complex  substances  are  pre- 
cipitated in  the  absence  of  chloride.  What  is  the  valence  of  the  OH  radical? 

Of  the  NH2  radical? 

(e)  H2S  in-  acid  solution. 

8.  Predict  the  action  of  dilute  HC1  or  HNO3  on  each  of  the  precipitates  just 
obtained  (except  HgS),  and  test  your  answers  by  experiments. 

[59] 


9.  Experiment.     Collect    some    HgS    on    a    filter    and    wash    it    with    water. 
Transfer  part  of  the  precipitate  to  a  porcelain  dish,  add  2  TV  HNO3  boil  the 
mixture  for  about  a  minute,  and  filter.     Test  the  filtrate  for  Hg*"1".     Repeat  the 
experiment  with  hot  2  N  HNO3  to  which  some  HC1  has  been  added.     From  the 
fact  that  HgQ2  is  a  weak  salt  suggest  an  explanation  for  the  result  of  this 
experiment.     Repeat  the  experiment  with  bromine  solution  instead  of  HNO3. 

10.  Other  reactions  of  mercurous  ion.     Experiment.     Treat  a  dilute  solution 
of  HgNO3  with  each  of  the  following  reagents : 

(a)  H2SO4  or  a  soluble  sulfate. 

(b)  Na2C03. 

(c)  NaOH.     The  precipitate  is  Hg2O,  which,  on  heating,  decomposes  into  Hg 
and  HgO. 

(d)  NH4OH.     The  precipitate  is  a  mixture  of  metallic  mercury  and  an  am- 
monobasic  mercuric  salt;  see  7  (d). 

(e)  H2S  in  acid  solution.    The  precipitate  is  a  mixture  of  Hg  and  HgS. 

11.  From  the  results   of  your  experiments  in   Paragraphs  6,  8,  and  9,   give 
methods  of  dissolving  each  of  the  precipitates  obtained  in  10.     Try  experiments 
whenever  you  are  doubtful  of  your  method. 

12.  Table  of  reactions  for  mercurous  ion  and  mercuric  ion.       Summarize  in 
a  table  for  future  reference  the  reactions  of  Hg*  and  Hg++  with  the  following 
reagents:  HC1;  a  strong  oxidizing  agent  as  Br2  a  strong  reducing  agent  as 
SnCl2   (in  small  amount  and  in  excess),  and  the  reagents  used  in  Paragraphs  7 
and  10.     Below  this  table  classify  the  precipitates  according  to  reagents  used  in 
dissolving  them  and  point  out  the  considerations  that  led  to  the  choice  of  each 
reagent.     Be  sure  that  you  can  write  the  equation  for  each  reaction. 

13.  Problems.  Try  experiments  when  necessary. 

(/)  Complete  the  following  equations,  and  indicate  what  you  would  observe 
in  the  corresponding  experiments: 

Sn+++  HgCl2  (excess)  = 
Sn++  (excess)  +  HgCl2  = 

(2)  What  reactions  are  possible  between  a  sulfide  and  hydrochloric  acid? 
Between  a  sulfide  and  nitric  acid?  In  connection  with  your  answers  discuss 
the  following  experimental  facts;  CuS  dissolves  readily  in  hot  2  N  HNO2  but 
not  in  hot  2  N  HC1. 

(j)  Compare  the  action  of  hot  2  Ar  HNO3  on  CuS  (Assignment  37  Paragraph 
10)  and  HgS,  and  state  what  conclusion  can  be  drawn  with  regard  to  the  relative 
solubilities  of  these  two  sulfides  in  water.  Outline  a  method  of  preparing  pure 
HgCl2  and  pure  CuSO4  from  a  mixture  of  the  freshly  precipitated  sulfides. 

(4)   Suggest  methods  of  making  the  following  preparations: 

(a)  HgO  fromHg(N03)2; 

(b)  Hg(NO3)2  from  HgSO4; 

(c)  Hg(N03)2  fromHgN03; 

(d)  HgNO3  from  Hg(NO3)2. 


Note.  It  is  recommended  that  Assignment  52  on  Qualitative  Analysis,  follow 
Assignment  44. 


60 


ASSIGNMENT  45 
LEAD  ION,  CHROMATE  ION 


1.  In  the  present  Assignment  and  the  following  ones  we  shall  continue  the 
study  of  important  common  elements.     Only  very  brief   laboratory  directions 
will  be  given;  each  student  is  expected  to  plan  the  details  of  the  experiments, 
and  is  recommended  to  try  additional  experiments  which  are  suggested  by  the 
lectures  or  laboratory  work. 

2.  Chromate    ion,   CrO4~~,   and   bichromate    ion,   Cr,O7~~.      Note  the  colors 
.  of   solutions  of  sodium    (or  potassium)    chromate  and  bichromate.     What  is 

the  color  of  CrO4~~?  Of  Cr2Or~~?  What  is  the  valence  of  chromium  in 
these  ions?  Experiment.  Add  dilute  HNO3  or  H2SO4  to  a  small  quantity 
of  a  chromate  solution.  Add  dilute  NaOH  or  NH4OH  to  a  bichromate  so- 
lution; acidify  the  solution  obtained.  Note.  Many  chromates  are  difficultly 
soluble  in  water,  but  dissolve  in  HNO3. 

3.  Reactions    of    lead    ion,  Pb++.       Experiment.       Treat    small    portions    of 
0.1   A7  Pb(NO3)2  with  each  of  the  reagents    listed    in    Assignment    38     (Par- 
agraph 5),  with  H2SO4,  and  with  a  soluble  chromate. 

4.  Methods  of.  dissolving  difficulty  soluble  lead  salts.     State  which  of  the  pre- 
cipitates obtained  in  (3)  you  would  expect  to  dissolve  in  0.2  N  HNO3.     Experi- 
ment.    Prepare  some  PbS,  collect  it  on  a  filter,  and  test  portions  of  it  with  hot 
2  N  HC1,  and  with  hot  2  N  HNO3.     Experiment.     Collect  some  PbCl,  on  a 
filter,  pour  cold  water  over  it  and  test  the  filtrate  for  Pb*+,  e.  g.  with  H2S; 
then   pour  hot   water  over   PbCl2   on   a   filter  and   test   the   filtrate   for   Pb++. 
Experiment.     Collect   some  PbCrO4  on   a  filter,  wash  thoroughly  with   water 
to  remove  any  soluble  chromate,  and  determine  how  readily  it  dissolves  in  hot 
dilute  HNO3.    What  ions  are  present  in  the  solutions  obtained?    Experiment. 
Collect  some  PbSO4  on  a  filter,  transfer  some  of  it  to  a  test  tube,  and  pass  in 
H2S.     Which  is  the  less  soluble,  PbSO4  or  PbS?     Experiment.     Treat  some 
PbSO4  with  Na2CO3  solution,  as  directed  in  Assignment  52,  11  (5). 

5.  Relative  delicacy  of  tests  for  Pb++.  Experiment.     Prepare  50  or  100  cc. 
portions  of  0.01   N,  0.001  N,  0.0001  N  Pb(NO3)2.  Test  5  or  10  cc.  of  each 
solution  with  the   reagents  used  in    (3).  Specify  the  two   most  delicate  tests 
Name  the  two  least  soluble  salts  of  lead. 

6.  Lead  dioxide,  PbO2.     Experiment.     Treat  a  small  portion  of   PbO2  with 
hot  6  N  HC1.    What  gas  is  evolved?    Repeat  the  experiment  with  HNO3  instead 
of  HC1. 

7.  Problems.     (/)   Outline  methods  for  making  the  following  transformations: 

(a)  PbSO4  to  PbS.  (c)  PbCO3  to  PbCrO4. 

(b)  PbS  to  PbSO4.  (d)  PbO2  to  Pb(NO3)2. 

(2)   Name  three  amphoteric  hydroxides,   (cf.  Assignment  35).     What  would 
be  observed  on  gradually  acidifying  a  solution  of  sodium  plumbite? 


ASSIGNMENT  46 
STANNOUS  AND  STANNIC  IONS.    AMPHOTERIC  SULFIDES 


Reference.    Hildebrand,  pages  197-199. 

1.  Ammonium  polysulfide  solution.  Experiment.  Prepare  some  colorless 
ammonium  sulfide  solution  by  passing  H2S  into  NH.tOH  solution.  Divide  the 
solution  into  three  parts.  To  one  portion  add  a  very  small  amount  of  powdered 

[61  1 


sulfur.  Place  another  portion  on  a  watch-glass  to  expose  it  to  the  action  of  the 
oxygen  of  the  air.  Acidify  the  third  portion  with  dilute  HC1.  Experiment. 
Acidify  small  quantities  of  the  laboratory  solutions  of  ammonium  polysulfide 
(yellow  ammonium  sulfide),  and  of  ammonium  sulfide.  Note.  On  account  of 
hydrolysis  the  (NH4)2S  solution  is  really  NH4OH  +  NH4SH. 

2.  Reactions    of   Sn++  and  Sn++++.    .  Experiment.       Prepare    0.1    N    solutions 
of  stannous  chloride  and  of  stannic  chloride,  and  determine  the  action  of  the 
following  reagents  on  each  of  these  solutions : 

(a)  H2S  in  cold  0.3  N  HC1  solution.    These  tests  are  delicate  and  characteristic. 

(b)  H2S  in  hot  2  N  HC1.    If  there  is  a  precipitate,  filter,  add  water  to  the  filtrate 
and  test  it  for  tin. 

(c)  NH4OH  in  small  amount  and  in  excess. 

(d)  NaOH  in  small  amount  and  in  excess.  If  there  is  a  precipitate  after  adding 
NaOH  in  excess,  filter,  acidify  the  filtrate  with  HC1  or  H2SO4  and  test  it  for 
tin. 

0)    (NH4)2S  after  NH4OH  has  been  added. 

(/)   Ammonium  polysulfide  after  NH4OH  has  been  added.    Pour  through  a  filter 

to  remove  any  precipitate  that  may  remain  undissolved,  and  acidify  with 

dilute  HC1. 

(g)  A  soluble  carbonate.     Filter,  wash  the  precipitate  and  test  it  for  carbonate. 
(h)  A  good  oxidizing  agent  in  dilute  acid  solution,  such  as  bromine. 
(i  )   A  good  reducing  agent  in  dilute  acid  solution,  as  zinc. 
(/  )   HgCl2  solution  in  small  amount  and  in  excess. 

3.  Action  of  concentrated  HNO^  on  Sn.       Experiment.      Treat  a  small  piece 
of   tin   with   concentrated   HNO3   in   a   casserole.     Evaporate   the   mixture   just 
to  dryness:  to  avoid  heating  the  residue  strongly,  evaporate  the   last  portions 
of  the  liquid  by  moving  the  casserole  back  and  forth  through  a  small  flame. 
Warm  the  residue  with  dilute  HNO3,  filter,  and  test  the  filtrate  for  tin.    Note. 
Stannic  acid,   or  hydrated   stannic   oxide,   when  prepared  in   this   way  by   the 
action  of  concentrated  HNO3  on  Sn  or  on  a  compound  of  tin,  is  difficultly 
soluble  in  HNO3.     Experiment.     Suggest  and  try  methods  of  dissolving  this 
substance. 

4.  Problems,     (i)     If  a  solution  were  known  to  be  either  SnCL  or  SnClt  how 
would  you  identify  it? 

(2)   How  would  you  test  for  SnQ2  in  a  solution  of  SnCl4? 

(j)  What  happens  when  HC1  is  added  gradually  to  a  solution  of  (a) 
sodium  stannite,  (b)  sodium  stannate,  (c)  sodium  stannite  after  it  has  been 
treated  with  a  powerful  oxidizing  agent,  as  sodium  peroxide,  Na2O2,  or  hypo- 
chlorite,  (d)  sodium  sulfostannate? 

(4)  How  would  you  prepare  (a)  ammonium  polysulfide,  (b)  sodium  sulfo- 
stannate from  SnS2,  (c)  sodium  sulfostannate  from  SnS? 


ASSIGNMENT  47 
IONS  OF  ALUMINUM  AND  OF  CHROMIUM.  PEROXIDES 


1.  Write  the  formulas  of  aluminum  chloride,  nitrate  and  sulfate;  chromium 
chloride,  nitrate  and  sulfate;  alum  and  chrom  alum.    These  salts  are  all  readily 
soluble  in  water  and  their  solutions  react  acid  to  litmus. 

2.  Hydroxides.      Experiment.       To    very    dilute    solutions    of    an    aluminum 
salt  and  of  a  chromium  salt  add  NH4OH  a  few  drops  at  a  time  until,  after 
shaking,  the  odor  of  ammonia  is  just  perceptible.     Warm  the  mixture.    These 
are  delicate  tests  for  Al+++  and  Cr^*. 

T621 


3.  When    a   solution   of   an   aluminum,    or   chromium,   salt   is   treated   with    a 
soluble  carbonate  or  sulfide,  the  hydroxide  of  aluminum,  or  chromium,  is  pre- 
cipitated.    Explain. 

4.  Aluminate   ion   and  chromite   ion.   A1O2~    and    CrCX"    (or    H.AICX')    and 
H,CrO3~).     Experiment.     To  dilute  solutions  of  an  aluminum  salt  and  of  a 
chromium   salt  add   a   few   drops   of   NaOH   solution.     To  each   mixture   add 
more  NaOH  and  shake  after  each  addition.     Filter  if  a  precipitate   remains, 
add  some  solid  ammonium  salt,  and  heat  the  mixture. 

5.  Ions  of  Chromium.     State  which  ions  of  chromium  can  exist  only  in  an 
acid    solution   and   which   in   an    alkaline    solution.      What    is    the  valence    of 
chromium  in  each  of  these  ions?     Suggest  a  method  of  changing  bichromate 
ion  to  chromic  ion. 

6.  Reduction  of  Cr2O7~~  and  of  CrO.,~~.       Experiment.       Test    the    action 
of  H2S  on  very  dilute  solutions  of  a  bichromate  in  the  presence  of   sulfuric 
acid,  heat  the  mixture  to  boiling,  filter,  and  wash  the  precipitate.     Make  the 
filtrate  alkaline  with  NH4OH.     Repeat  the  experiment  with  H2S  and  a  very 
dilute  solution  of  a  chromate  in  the  presence  of  NaOH,  heat,  filter  and  wash 
the  precipitate.    Acidify  the  filtrate.    Experiment.     Determine  the  action  of  hot 
concentrated  HC1  on  Cr2O7~~,  and  of  Cl~  on  CrO4~~  in  alkaline  solution. 

7.  Oxidation    of    chromite    ion.       Experiment.        Prepare    a    small    quantity 
of  a  chromite  solution.     To  a  portion  of  this  add  bromine  solution;  boil  the 
solution  to  expel  any  bromine  that  may  remain.     Question.     How  would  you 
prepare    a   hypochlorite   solution?      Write    the    equation    for   the   oxidation    of 
CrO2~  by  C1O~.     Experiment.     Boil  a  small  quantity  of  a  chromic  salt  solution 
with  concentrated  HNO3,  or  with  aqua  regia. 

8.  Peroxides.     Sodium  peroxide  decomposes  rapidly,  even  in  a  cold  solution. 

NiO         HO  =  2NaOH 


Hydrogen  peroxide  in  a  cold,  dilute  acid  solution  decomposes  only  very 
slowly.  Na2O2  is  a  very  powerful  and  useful  oxidizing  agent.  It  reacts  with 
chromic  hydroxide  suspended  in  water,  as  follows  : 

2Cr(OH)3  +  3Na2O2  =  2CrO4"  +  6Na+  +  2OH~  +2H,O. 
Experiment.     Test  the  action  of  Na2O2  on  CrO2~.     The  evolution  of  oxygen 
observed    in    this    experiment    is    due    to    the    decomposition    of    some    of    the 
peroxide. 

9.  In  acid  solution  H2O2  may  be  reduced  by  a  powerful  reducing  agent,  as 
HoS;  but  it  may  also  be  oxidized  by  a  powerful  oxidizing  agent,  as  Cr2O7~~. 

Cr2O7"  +  3H2O2  -f  8H+  =  2  Cr+++  +3O2  +  7H2O. 

10.  Problems.     (/)     Complete  the  following  equations  and  state  what  would 
be  observed  in  the  corresponding  experiments: 

Cr(OH)3    +    H+    (excess)  =  .................................................... 

CrO2-  +  H+  (excess)  =  .................................................... 

CrO4-  +  H+  (excess)  =  .................................................... 

Cr20r~  +  H2S03  +  H+  =  Cr+"  ............................................ 

H2O2  +  MnO2  +  H+  =  Mn"  ............................................ 

H202  +  I-  +  H+  =  I,  ................................................. 

(2)    State  how  you  would  prepare  Cr2Or~~  from  Cr+++,  Cr(OH)3  from  PbCrO4. 

(j)  Concentrated  NH4OH  solution  has  a  slight  solvent  action  on  Al(OH)g 
and  Cr(OH)3,  but  these  hydroxides  can  be  reprecipitated  by  adding  an  ammon- 
ium salt  and  warming  the  mixture.  Explain. 

(4)  List  in  a  vertical  column  all  the  amphoteric  hydroxides  you  have  studied, 
and  write  in  two  more  columns  the  formulas  of  the  corresponding  ions  which 

[63] 


are  formed  when  the  hydroxide  is  dissolved  in.  excess  of  a  solution  of   (a)   a 
strong  acid,  and  (b)  a  strong  base. 

(5)  Examine  the  position  in  the  Periodic  System  (Hildebrand,  page  257)  of 
each  of  the  elements  considered  in  the  preceding  problem,  and  write  a  brief 
note  on  the  gradations  of  properties  in  each  family  in  which  one  of  these 
elements  occurs.  Name  additional  elements  which  might  be  expected  to  have 
amphoteric  hydroxides. 


64 


SECTION  V 
QUALITATIVE  ANALYSIS 


ASSIGNMENT  51 

THE  DEVELOPMENT  OF  A   SCHEME  OF  ANALYSIS  FOR  A  LIMITED  NUMBER  OF 

POSITIVE  IONS 


(To  follow  Assignment  38) 

1.  In  Assignment  38   the  chemistry  of   the  positive  ions  already   considered 
was  summarized,  and  the  reactions  of  these  ions  with  various  reagents  tabulated. 
We   shall   now   consider   how   this   information  may   be   used   in   analyzing   an 
unknown  solution  for  these  ions. 

2.  A  possible  method  of   analysis  would  be  to  devise  a  separate  procedure 
for  each  ion,  then,  if  seven  ions  were  to  be  considered,  seven  separate  portions 
of  the  original  solution  (or  substance)  would  be  taken,  and  each  treated  by  an 
experiment   or  series   of   experiments   devised   to   show   whether   the   particular 
substance  is  present  or  absent  in  the  unknown.     Such  tests  were  used  in  Assign- 
ment  31    for   Cl~,   SO4~~,   and   NO3~;   in   these   cases   provision   was   made   for 
neutralizing  OH~,  because  the  presence  of  OH"  would  have  interfered  with  or 
spoiled  the  tests.     It  is  to  be  noted  also  that  the  test  used  for  Cl~  is  not  specific 
for  Cl~,  since  any  other  silver  salt  which  is  insoluble  in  dilute  HNO3,  as  silver 
iodide,   would  also  be  precipitated.     It  is   evident   that   the  application  of  this 
method  (of  testing  separate  portions  of  the  unknown  for  a  particular  substance) 
becomes  increasingly  difficult  as  the  number  of  substances  which  may  be  present 
is  increased.     Also  a  lengthy  procedure  is  often  necessary,  especially  in  testing 
for  certain  positive  ions  when  ions  of  several  other  metals  are  present. 

3.  Accordingly,  in  systematic  qualitative  analysis  the  method  has  been  adopted 
of  using  only  one  portion  of  the  unknown  solution  in  testing  for  metals,  and 
of  preparing  from  it  a  pure  compound  of  each  metal.     Such  a  process  serves 
to  identify  the  positive  constituents  and  also  enables  a  comparison  of  the  amounts 
of  the  various  metals  to  be  easily  made.  The  reagents  ordinarily  used  in  this 
process  are  acids,  NH4OH  and  NH4  salts;  and  therefore  separate  portions  of 
the  unknown  are  used  in  testing  for  H+,  NH4+,  and  the  various  negative  ions. 
(The  development  of  a  systematic  procedure  for  the  detection  of  negative  con- 
stituents will  not  be  considered  in  this  course.) 

4:  As  an  illustration  we  shall  discuss  the  preparation  of  pure  salts  of  zinc  and 
calcium  from  an  acid  solution  containing  Zn++  and  Ca++.  The  Tables  in  Assign- 
ment 38  show  the  following  differences  which  may  be  utilized  in  devising  methods 
of  separation : 

(/)  CaS  is  soluble  in  water  and  ZnS  is  difficultly  soluble.  The  latter  can 
be  precipitated,  after  the  solution  has  been  made  alkaline  with  NH4OH,  by  add- 
ing (NH4)2S  solution  drop  by  drop,  or  better  by  passing  in  H2S  gas.  When 
the  ZnS  has  been  completely  precipitated,  and  removed  by  filtration,  calcium 
can  be  tested  for  in  the  filtrate  and  any  desired  salt  of  calcium  can  be  pre- 
pared. Experiment.  Test  this  method  by  the  Procedure  given  below,  with  an 
acid  solution  which  contains  both  Zn++  and  Ca++,  prepared  for  example  by  adding 
to  100  cc.  water  in  a  flask  2  or  3  drops  of  ZnSO4,  CaCL  and  HC1  solutions. 
Repeat,  omitting  the  calcium  salt,  in  order  to  determine  if  a  satisfactory  blank 
is  obtained  in  the  calcium  test.  Procedure.  (Be  sure  that  you  understand  the 
reason  for  each  detail  in  the  directions.)  Add  NH4OH,  a  few  drops  at  a  time, 
until  the  mixture  after  shaking  has  a  distinct  odor  of  NH3,  pass  in  H2S,  shake 

[651 


the  mixture  thoroughly,  and  cautiously  note  the  odor  to  determine  if  H2S  has 
been  added  in  excess;  pass  in  more  H2S  if  necessary.  (If  a  large  precipitate 
is  obtained,  test  with  litmus  and  if  the  solution  has  become  acid  add  more 
NH4OH;  again  pass  in  H2S  until  it  is  present  in  excess.)  Warm  the  mixture 
to  50  or  60°,  and  filter.  Test  the  clear  filtrate  for  calcium  by  adding  (NH4)2CO3 
solution,  and  heating  the  mixture  almost  to  boiling.  Note.  If  in  the  analysis 
of  an  unknown  solution  no  precipitate  is  observed  in  the  test  for  Zn++  or  Ca++, 
let  the  warm  solution  stand  for  several  minutes  before  concluding  that  the  ion 
is  absent. 

(2)  Zn++  forms  a  complex  ion  with  ammonia  and  Ca++  does  not.  Both 
carbonates  are  difficultly  soluble  in  water,  but  ZnCO3,  unlike  ZnS,  dissolves  in 
NH4OH.  Accordingly  CaCO3  can  be  precipitated  by  treating  the  mixture  with 
NH4OH  in  excess  and  (NH4)2CO3.  Write  out  a  Procedure  (cf.  preceding  Par- 
agraph) and  note  the  reason  for  each  of  the  experimental  details. 

(5)  Zn++  is  converted  into  zincate  ion  by  excess  of  a  strong  base.  Pure 
NaOH  solution  would  precipitate  some  of  the  calcium  as  Ca(OH)2  but 
Na2CO3  must  also  be  added  to  make  a  satisfactory  separation.  The  laboratory 
NaOH  solution  contains  some  CO3~~  but  it  is  better  to  add  Na2CO3. 

5.  Plan   three  methods   of   separating  silver  and   calcium,   i.   e.   of   preparing 
pure  salts  of  silver  and  of  calcium  from  a  mixture  of  their  nitrates.     Write 
out  each  method,  stating  why  you  think  it  will  work;   and  show  your  note- 
book to  the  instructor. 

6.  By    combining  your   methods   of   separation,    Paragraphs   4   and   5    devise 
a  "scheme  of  analysis"  for  these  three  positive  constituents,  silver,  zinc  and 
calcium.    Test  your  scheme  by  experiments  with  very  small  quantities  of  solu- 
tions of   (a)   AgNO3  alone,   (b)   ZnSO4  alone,  and   (c)    CaQ2  alone.     In  each 
experiment  with  a  pure  salt  a  satisfactory  test  should  be  obtained  for  one  metal 
and  blanks  in  the  tests  for  the  other  two. 

7.  Experiment.     Predict  what  will  happen  when  a  small  quantity  of  CuSO4 
solution  is  treated  by  your  scheme  of  analysis,  Paragraph  6,  and  test  your  answer 
by  an   experiment.   Devise  a  method  of  including  copper  in  your  scheme,  and 
test  your  method  by  experiments. 

8.  Assume   that   a   sodium    salt   is   present   in    solution   with   salts   of    silver, 
copper,  zinc  and  calcium,  and  consider  how  all  the  sodium  originally  present, 
and  no  more,  could  be  recovered  as  a  pure  salt.   In  other  words,  include  sodium 
in  your  scheme.     Show  your  scheme  of  analysis  for  the  five  positive  constit- 
uents  to   the   instructor.       Question.       With   which  of   the   five   metals   would 
potassium  have  been  found  if  it  had  been  present  in  the  original  solution?    -0. 

9.  When  an  unknown  solution  has  been  analyzed  in  this  way  and  a  result 
is  inconclusive,  owing  for  example  to  the  small  size  of  the  precipitate,  a  charac- 
teristic  confirmatory    test   is   made    whenever    possible,    such    as    the    dissolv- 
ing of  AgCl  in  NH4OH  and  the  reprecipitation  with  excess  HNO3,  the  flame 
test  for  calcium,  etc. 

10.  When   the   unknowns  are    solid   salts   an   important   part   of   the   work   is 
the  preparation  of  the  solution  for  analysis.  Preliminary  experiments  with  small 
quantities  are  performed  to  determine  a  method  of  dissolving  the  unknown,  and 
conclusions  with  regard  to  its  nature  can  often  be  drawn  from  the  results  of 
these  experiments.     The  salts  given  out  for  analysis  at  this  time  will  dissolve 
either   in   water,    in   dilute   acids   or   in   concentrated   acids.    .  A   discussion   of 
methods  of  dissolving  difficultly  soluble  salts  will  be  given  in  Assignment  52. 

11.  Review    your    notes    in    the    previous    Assignments    on    the    tests    for 
Na+,  K+,  NH4+,  C1-,   SO4",  CCX"  and  NO3-.  Devise  a  method  of  testing  for 

[661 


S"  based  on  the  formation  of  H2S  gas,  and  determine  whether  your  test  will 
apply  in  the  case  of  a  very  difficultly  soluble  sulfide,  as  CuS. 

12.  Analyses   1,   2  and  3.      (Note.     Do  not  begin   these  analyses  until  the 
instructor  has  approved  your  scheme  of  analysis   for  the  positive  constituents, 
Paragraphs  6  to  8. )     Test  for  the  positive  and  negative  constituents : 

H,  Ag,  Cu,  Zn,  Ca,  Na,  K,   NH4 
OH,  Cl,  N03,  S04,  C03,  S. 

Record  at  once  each  experiment  and  observation  and  write  down  any  con- 
clusion that  has  bearing  on  the  identification.  Try  to  distinguish  between  large 
amounts,  small  amounts,  and  traces  of  the  different  constituents.  When  you 
report  the  result  of  an  analysis  to  your  laboratory  instructor  show  him  your 
notebook  record. 

13.  Problems.     (/)     What  negative  ions  need  not  be  tested  for  in  solutions 
which  contain  the.  following  positive  ions:    (a)    H+.    (b)    Ca+*,    (c)    Zn++,    (d) 
Cu++,  (e)  Ag*? 

(2)   Point  out  the  errors  in  the  following  reports: 

(a)  A  solution  known  to   contain  Ag+  was   reported  to   contain   SO4~ ' 
because  BaCL  gave  a  white  precipitate  which  did  not  dissolve  in 
HN03. 

(b)  A  pure  sodium  salt  which  dissolved  in  water  to   give  an  alkaline 
solution  was  reported  to  be  a  sulfate  because  a  white  precipitate 
was  obtained  on  the  addition  of  Ba(NO3)2- 

(c)  A  pure  salt  which  did  not  dissolve  appreciably  in  water  was  reported 
to  be  zinc  nitrate. 


ASSIGNMENT  52. 

THE  STANDARD  SCHEME  OF  ANALYSIS.     METHODS  OF  DISSOLVING  DIFFICULTLY 

SOLUBLE  SUBSTANCES 


(To  follow  Assignment  44) 

1.  In  Assignment  51  a  number  of  schemes  were  considered  for  the  systematic 
analysis  of  solutions  containing  Ag+,  Cu"+,  Zn++,Ca  ++  and  Na+  (and  K+).  One  of 
the  most  satisfactory  schemes,  and  the  one  used  in  practically  all  text  books, 
is  the  following: 

To  the  solution  add  HC1 


Precipitate 
AgCl 


Filtrate  (concentration  of  H+,  0.3  Ar)  :  Saturate  with  H2S  gas 


Precipitate 
CuS 


Filtrate :  Boil  to  expel  H0S,  *  add  NH4OH  and 
(NH4)2S. 


Precipitate 
ZnS 


Filtrate:  Add  (NH4),CO3  and  warm 


Precipitate    Filtrate :  Evaporate  to  dry- 
CaCO2  ness, 

[eat  to  expel  NH4  salts. 
Residue:    Na     (and    K) 
salts. 

*  Note.  It  is  really  not  necessary  to  expel  the  H2S ;  but  when*  additional  positive  con- 
stituents are  included  in  the  scheme,  and  several  substances  may  be  precipitated  at  this 
point,  the  observation  of  the  effect  of  adding  NH4OH  in  the  absence  of  sulfide  often  assists 
in  identifying  the  positive  constituents. 

[67] 


2.  Iron  and  mercury  in  the  scheme  of  analysis.     From  the  results  of  your 
experiments  on  the  reactions  of  the  ions  of  iron  and  mercury,  Assignments  43 
and  44,  predict  what  would  be  observed  if  a  solution  containing  each  of  the 
following    ions     was     treated     by     the     scheme     of     analysis,     Paragraph     1 : 
Fe++,  Fe+++,  Hg+,  Hg++.     Experiment.     Verify  your  answers  by  experiments  with 
dilute  solutions  of  each  of  these  ions.     Be  sure  that  the  blank  tests  for  the  other 
positive  constituents,  e.  g.  Ca++,  are  negative.     Question.  What  mistake  has  been 
made  if  in  the  experiment  with  a  mercuric  salt  a  black  precipitate  is  obtained 
with  sulfide  in  the  NH4OH  solution? 

3.  Separation  of  ZnS  and  FeS.     Plan  a  method  of  dissolving  a  mixture  of 
ZnS  and  FeS  and  of  obtaining  pure  Fe(OH)3  and  ZnS  from  the  resulting  solution 
(which  contains  Zn++  and  Fe++).     Experiment.     Test  your  method  with  small 
amounts  of  (a)  ZnS,  (b)  FeS,  and  (c)  a  mixture  of  ZnS  and  FeS.  See  Assign- 
ment 51  Paragraph  4  (/)  on  the  method  of  percipitating  ZnS. 

Separation  of  CuS  and  HgS.  Experiment.  Prepare  a  small  amount,  of  (a) 
CuS,  and  (b)  a  mixture  of  CuS  and  HgS.  Treat  each  of  these  by  the  procedure 
devised  in  Problem  (j),  Assignment  44.  Questions.  Why  is  the  sulfur  residue 
black  when  CuS  alone  is  used,  (cf:  Assignment  37,  Paragraph  10)?  Is  a 
confirmatory  test  for  mercuric  mercury  necessary?  For  such  a  test  see  Assign- 
ment 44,  Paragraph  3. 

Separation  of  AgCl  and  HgCl.  Devise  and  confirm  by  experiment  a  method 
of  determining  whether  a  chloride  precipitate  is  either  AgCl  or  HgCl.  If  a 
large  amount  of  HgCl  is  found  and  only  a  small  amount  of  AgCl  is  present, 
the  method  based  on  the  use  of  NH4OH  solution  fails  on  account  of  the  reduction 
of  Ag+  by  the  metallic  mercury.  Devise  a  method  of  detecting  silver  in  the 
presence  of  a  large  amount  of  HgCl,  based  on  the  treatment  of  the  mixture 
obtained  on  the  addition  of  NH4OH,  with  an  oxidizing  agent  such  as  bromine 
solution.  Experiment.  Try  your  method. 

4.  The   other    common   positive   elements    are   grouped   as    follows    with    the 
elements  already  considered. 

Division  of  Metals  into   Groups  for  Qualitative  Analysis. 

Characteristic  Reagents    for     Precipita-     Substances    Precipitated. 

Metal  tion. 

Name  of  Group. 


I.       Silver 

Silver  Group. 

II.       Copper 

Copper  Group. 


[II.       Zinc 

Zinc   Group. 


IV.       Calcium 

Calcium  Group. 


HC1    or    NH4C1    and   an     AgCl,   HgCl,   PbCl2   par- 
tially,  BiOCl   partially. 


acid. 
H»S  in  0.3  N  acid. 


NH4OH  and  (NH4)2S 
or  NH4OH  in  excess 
and  H,S. 


CuS,  PbS,  HgS,  CdS, 
Bi2S3,  Sulfides  of  As, 
Sb,  Sn. 

ZnS,  A1(OH)3,  Cr(OH)s 
FeS,  Fe2S3,  MnS,NiS, 
CoS. 


(NH4)2CO3andNH4OH.     CaCO3,  BaCO3,  SrCO3,  a 

compound      of      rnag- 
nesium  when  alcohol  is 
present, 
none. 


V.       Sodium 

Sodium  Group.  , 

5.  The    further   separation    of   the   elements    has   already    been   illustrated   in 
Paragraphs  2  and  3,  but  only  for  two  elements  in  each  of  three  groups.    The  task 

F681 


of  developing  a  scheme  of  analysis  becomes  increasingly  difficult  as  the  number 
of  elements  is  increased, — mainly  on  account  of  the  large  number  of  time-con- 
suming experiments  which  must  be  made  with  known  mixtures  in  order  to 
prove  that  the  proposed  methods  are  satisfactory.  Therefore  if  you  ever  have 
to  make  an  accurate  qualitative  analysis  of  a  complicated  mixture  it  will  be 
necessary  to  consult  a  standard  text-book,  such  as  A.  A.  Noyes'  Qualitative 
Chemical  Analysis,  and  to  follow  carefully  the  directions  there  given.  Before 
attempting  to  make  an  analysis  with  the  aid  of  a  book,  work  through  the  pro- 
cedures with  known  solutions  and  study  the  chemistry  involved. 

6.  The   present    course,    however,    is    not   primarily    a    course    in    Qualitative 
Analysis,  and  a  text-book  will  not  be  necessary.     In  the  time  which  remains  at 
our  disposal  only  a  limited  number  of  additional  elements  can  be  studied  in  the 
laboratory.     These  will  be  investigated  by  the  same  methods  as  before,  and 
from  the  tables  of  reactions  it  will  be  possible  to  develop  satisfactory  methods 
for  the  analysis  of  simple  mixtures. 

7 .  Analyses  Nos.  4,  5  and  6.     In  addition  to  the  positive  and  negative  con- 
stituents  considered   in  Assignment   51,  test   for  iron,   mercurous   mercury   and 
mercuric    mercury.         Do    not    begin    these    analyses    until    the    instructor   has 
approved  your  methods  of  analysis,  Paragraphs  2  and  3.     Also  study  carefully 
the  following  paragraphs. 

8.  Analysis  of  an  unknown  solid.     Record  the  appearance  of  the  solid.     Treat 
a  small  portion  with  water  and  heat  the  mixture  in  a  porcelain  dish  or  beaker 
if  the  solid  does  not  dissolve  readily.     Always  test  the  solution  with  litmus  paper. 
If  the  solution  is  neutral,  determine  if  anything  has  dissolved,  either  by  slowly 
evaporating  a  few  drops  of  the  clear  solution  on  a  watch-glass,  or  by  quickly 
evaporating  a  larger  quantity  of  the  solution  in  a  porcelain  dish. 

9.  If  the  solid  does  not  dissolve  readily  in  water,  in  order   to  find  a  satis- 
factory method  of  dissolving  it,  try  preliminary  experiments,  such  as  are  sug- 
gested below,  with  small  quantities  of  the  powdered  substance.     The  problem 
is  much  simplified  if  the  solid  is  known  to  be  a  single  pure  substance;  for  in 
this  case,  when  a   reagent  is   found  to   react  with   some  of  the  solid,   the   re- 
mainder can  be  made  to  react  in  the  same  way  by  continued  treatment  with 
this   reagent. 

10.  When  a  method  of  dissolving  the   solid  has  been   found,   dissolve   about 
1   gram,  and  set  aside  about  two-thirds  of  the  solution  in  order  that  portions 
of  it  may  later  be  used  in  testing  for  NH4+  and  the  negative  constituents,  and  in 
making  confirmatory  tests.     Treat  the  remainder  by  the  scheme  of  analysis  for 
the  positive  constituents ;  use  beakers  or  flasks  in  this  work,  and  do  not  throw 
away  a  portion  of  a  solution  whenever  you  have  a  large  volume  of  it. 

11.  Methods  of  dissolving  a  difficultly  soluble,  pure  substance.     Observations 
in  these  preliminary  experiments  with  small  quantities  of  the  solid  may  be  of 
great  assistance  in  identifying  the  substance.     (In  the  following,  some  substances 
are  mentioned  which  will  not  be  encountered  in  Analyses  4,  5  and  6). 

(/)  Treatment  with  dilute  acid.  Use  N  HNO3,  boil  if  necessary.  Note 
the  odor  of  any  gas  evolved.  The  use  of  HC1  is  avoided  because  AgCl  or 
HgCl  might  be  formed.  Similarly  BaSO4  or  PbSO4  might  form  if  H2SO4 
were  used. 

(2)  Treatment   with   hot   concentrated   HC1.      If   chlorine   is   evolved   a   very 
powerful  oxidizing  agent  is  present,  such  as  MnO2  or  PbO2.     Concentrated  HC1 
dissolves  AgCl  and  HgCl  readily,  with  formation  of  complex  negative  ions,  but 
the  chlorides  precipitate  again  when  water  is  added  to  dilute  the  acid. 

(3)  Treatment    with    hot    concentrated    HNO3,    a   powerful    oxidizing    agent. 
The    rapid    formation    of    brown    gases    indicates    the    presence    of    a    reducing 

[691 


agent.     There  is  a  residue  of  sulfur  after  the  action  of  nitric  acid  on  a  sulfide, 
and  the  solution  may  contain  SO4~~. 

(4)  Treatment  with  either    (a)    a  mixture  of   concentrated  HC1  and  HNO3 
(aqua  regia)  or  (b)  bromine  water.      A  few  substances,  as  HgS,  are  dissolved 
more   readily  by  these  reagents  than   by  concentrated  HNO3    (which  is   really 
a  more  powerful  oxidizing  agent)  on  account  of  the  formation  of  weak  electrolytes 
or  complex  ions. 

Note.  Before  using  for  analysis  a  solution  containing  concentrated  acid 
evaporate  it  almost  to  dryness  to  remove  the  excess  of  acid,  and  add  water. 
Question.  What  happens  when  H2S  is  passed  into  concentrated  HNO3  or 
aqua  regia? 

(5)  Treatment    with    Na2CO3    solution    to    form    a   carbonate.      Use    a    con- 
centrated solution  of  Na2CO3  in  a  porcelain  dish,  cover  the  dish  with  a  watch- 
glass  and  boil  for  several  minutes.     Filter  and  retain  the  nitrate  to  test  for  the 
negative  constituent.     Treat  the  residue  with  dilute  HNO3,  filter  if  necessary, 
and  analyze  the   solution   for  the  positive   constituent.     Review   the   action   of 
Na.,CO3   solution  on   CaSO4,   Assignment   32;   and,    experiment,    try   the   above 
procedure  with  BaSO4. 

(6)  Special  tests.     For  example,   AgCl  dissolves  in  NH4OH   solution,  while 
HgCl  turns  black;  AgCl,  when  treated  with  water  and  H2S,  is  transformed  into 
the  black  Ag2S,  which  then  may  be  collected  on  a  filter  and  dissolved  in  con- 
centrated HI\fO3. 

12.  When  the  unknown  is  a  mixture  of  solid  salts,  the  same  method  of  pro- 
cedure may  be  followed  in  order  to  determine  a  method  of  dissolving  the  solid 
completely.  However,  it  is  often  convenient  to  analyze  separately  the  portions 
dissolved  in  the  different  operations.  For  example  the  mixture  might  consist  of 
one  salt  which  dissolves  readily  in  water  or  dilute  HNO3,  and  a  second  salt 
which  does  not.  If  this  method  is  followed,  care  should  be  taken  to  dissolve  the 
first  salt  completely  before  continuing  the  experiments  with  the  residue.  Note. 
Frequently  when  a  mixture  of  two  salts,  as  AgNO3  and  NaCl,  is  treated  with 
water  a  reaction  takes  place. 

L3.  Problems,  (i)  Briefly  outline  experiments  by  which  you  could  identify 
each  of  the  following:  AgCl,  HgNO3,  CuO,  HgS,  ZnCl2,  Fe(OH)3,  CaCO3. 

(2  In  order  to  be  able  to  interpret  quickly  observations  made  in  preliminary 
experiments  on  difficultly  soluble  substances,  Paragraph  11,  prepare  a  table, 
listing  in  a  vertical  row  all  the  difficultly  soluble  compounds,  studied  thus  far,  and 
in  a  horizontal  row  the  reagents,  (a)  dilute  HNO3,  (b)  hot  concentrated  HC1, 
(c)  hot  concentrated  HNO8,  (d)  bromine  water  (or  aqua  regia).  (c)  NH4OH 
solution :  If  the  compound  is  insoluble  in  a  given  reagent  mark  its  position  in 
the  table  with  a  cross,  if  soluble  write  in  the  formulas  of  the  substances  formed 
when  it  goes  into  solution. 


ASSIGNMENT  53 
LEAD,  TIN,  ALUMINUM,  CHROMIUM,  AND  BARIUM  IN  THE  SCHEME  OF  ANALYSIS 


(To  follow  Assignment  47) 
Reference.  Assignment  52. 


1.  Lead  in  the  scheme  of  analysis.  State  what  would  be  observed  if  a  solution 
of  pure  Pb(NO3)2  were  treated  by  the  scheme  of  analysis  which  you  used  in 
Assignment  52.  Sugest  a  method  of  determining  whether  a  white  precipitate 
obtained  with  HC1  contains  any  PbCL.  What  .would  be  the  result  of  treating 

[70] 


a  mixture  of  HgS,  CuS  and  PbS  with  hot  2  N  HNO3  (cf.  Assignment  52,  Par- 
agraph 3)  ?  Give  two  reactions  which  are  characteristic  of  Pb++  but  not  of  Cu++, 
and  one  reaction  which  is  characteristic  of  Cu++  but  not  of  Pb++.  Suggest  a 
method  of  separating  lead  from  copper  and  mercury,  i.  e.  of  including  lead  in  the 
scheme  of  analysis.  Experiment.  Test  your  method  with  known  solutions. 

2.  Tin  in  the  scheme  of  analysis.     State  what  would  be  observed  if  solutions 
of  (a)  stannous  tin,  (b)  stannic  tin  were  treated  by  the  scheme  of  analysis.     In 
which  analytical  group  would  stannous  and  stannic  tin  be  precipitated?     Suggest 
two  methods  of  separating  tin   from  copper,   mercury  and  lead.     Experiment. 
Try   one  of  your   methods   with   known   solutions.     Repeat   the   experiment  in 
Assignment  46,   Paragraph  2    (/),   using  small  amounts   of   dilute  solutions   of 
(a)  Cu(NO3)2,  (b)  SnCL,  and  (c)  a  mixture  of  Cu(N(X)2  and  SnCU. 

3.  Aluminum  and  Chromium  in  the  scheme  of  analysis.     State  what  would 
happen    if    separate  solutions  of    Al2(SO4)a,    Cr2(SO4)8    and    K2Cr.,O7    were 
treated  by  the  scheme  of  analysis.     In  which  analytical  group  are  aluminum  and 
chromium  precipitated,  and  what  are  the  formulas  of  the  substances  obtained? 
How  would  you  prepare  from  a  mixture  of  ZnS,  FeS,  A1(OH);<  and  Cr(OH)3 
an   alkaline  solution   containing   HZnO.r,   CrO2-   and  AlO.r?     What  would  be 
the   effect   of   adding   Na2O2   to   this   alkaline   solution?     Outline  a   method   of 
showing  the  presence  of  aluminum,  chromium  and  zinc  in  the  resulting  solution. 
Note.     The  alkaline  solution  which  contains  CrO4~~  must  not  be  acidfied  until 
the  peroxide  present  has  been  decomposed  by  boiling  the  solution.     Experiment. 
Test  your   method  with  known   solutions   containing  aluminum  and   chromium, 
and  chromium  alone.    Experiment.    Try  the  procedure  with  the  reagents.  NaOH 
and  Na2O2,  and  note  whether  a  satisfactory  blank  test  is  obtained  for  aluminum 
and  for  zinc. 

4.  Barium  in  the  scheme  of  analysis.     Explain  why  you  might  expect  barium 
to  be  found  with  calcium  in  the  scheme  of  analysis.     Look  up  in  a  table  of 
solubilities,  such  as  is  given  on  the  front  cover  page  of  Smith's  text-book,  the 
solubilities  in  mols  per  liter  of  the  hydroxides,  carbonates,  sulfates  and  chromates 
of  barium  and  calcium.     Arrange  these  salts  in  one  column  in  the  order  of  their 
solubilities  in   water.     Devise   two   methods   for  the   separation   of  calcium  and 
barium. 

5.  Procedure  for  the  separation  of  barium  and  calcium,  after  precipitation  as 
carbonates.    Experiment.    Test  the  following  procedure  with  a  freshly  prepared 
precipitate   of    (a)    BaCCX,   and    (b)    CaCO.v     To   dissolve   the   carbonate  pre- 
cipitate pour  repeatedly  through  the  filter  a   10  cc.  portion  of  hot  3  N  acetic 
acid,  and  use  a  little  more  acetic  acid  if  the  precipitate  does  not  dissolve  com- 
pletely.    Add   20   cc.  N   NH4Ac,   and  heat  the  solution  to  boiling  in   a  flask. 
Measure  out  15  cc.  A7"  Na2CrO4  and  add  it,  a  little  at  a  time,  heating  and  shak- 
ing the  mixture  after   each   addition.     Finally   heat   the  mixture   at   90°  -  100° 
for  one  or  two  minutes,  and  shake  it  at  the  same  time.     Filter,  even  if  you 
cannot  see  a  precipitate,  remove  the  filtrate  and  wash  the  precipitate  with  cold 
water.     (Pale  yellow  precipitate,  presence  of  barium.)  To  confirm  the  presence 
of  barium,  dissolve  the  precipitate  by  pouring  hot  HC1  through  the  filter,  evap- 
orating to  a  small  volume,  and  applying  the  flame  test.     To  test  for  calcium  in 
the   filtrate,-  make   the    solution   alkaline   with    NH4OH,    add    (NH4)2CO3,    or 
ammonium  oxalate,  and  heat  the  mixture  just  to  boiling.     Note.     The  separation 
depends  on  the  fact  that  BaCrO4  is  much  less  soluble  in  water  than  is  CaCrO4. 
If  the  BaCrO4  is  not  precipitated  slowly  in  a  hot  acid  solution,  as  directed,  the 
precipitate  will  be  so  finely  divided  that  it  will  run  through  the  filter. 

6.  Analyses    Nos.   7,  8,  9,  etc.      Test  for  all  the  positive  and  negative  con- 
stituents you   have   studied.     Verify   your   results   by   experiments   with   known 
substances.  ! 

[71  1 


7.  Problems.  (/)  Complete  the  table  of  the  behavior  of  difficultly  soluble 
substances  towards  various  reagents,  Assignment  52  Problem  (<?),  by  the 
addition  of  the -difficultly  soluble  substances  studied  in  Assignments  45-47  and  53. 

(2)   Point  out  the  errors  in  the  following  reports: 

(a)  Mercury  is  reported  to  be  present  because  a  black  residue  remained 
when  the  sulfides  of  the  copper  group  were  boiled  with  2  N  HNO3. 

(b)  Lead  is  reported  to  have  been  found  in  the  silver  group  but  not  in 
the  copper  group. 

(c)  In  the  analysis  referred  to  in   (b)  a  black  precipitate  was  observed 
in  testing  for  the  zinc  group,  although  no  test  for  iron  was  otbained. 
What  errors  in  the  method  of  analysis  would  produce  this  result? 

(d)  Aluminum  is  reported  to  be  present  in  the  zinc  group,  but  the  note- 
book shows  the  following  records :  on  precipitating  the  zinc  group 
a   greenish   precipitate   was   obtained   on    the   addition   of    NH4OH, 
later  a  yellow  solution  resulted  when  Na^CX  was  added ;  when  this 
solution   was  acidified   there   was   evolution  of   gas  and  the  yellow 
color   disappeared;   on  the   addition   of   NHtOH   a   precipitate   was 
formed. 


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